Praxis Chemistry Study Guide

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Chemistry Notes

Table of Contents

Glossary of Terms ….. ….. ….. ….. ….. ….. ….. ….. pg 2

Equations and Variables ….. ….. ….. ….. ….. ….. ….. ….. pg 13

Example Problems ….. ….. ….. ….. ….. ….. ….. ….. pg 14

Tables ….. ….. ….. ….. ….. ….. ….. ….. pg 19

Diagrams ….. ….. ….. ….. ….. ….. ….. ….. pg 22

Historical Figures ….. ….. ….. ….. ….. ….. ….. ….. pg 26

Applications ….. ….. ….. ….. ….. ….. ….. ….. pg 28

Nomenclature ….. ….. ….. ….. ….. ….. ….. ….. pg 29

Lab Techniques ….. ….. ….. ….. ….. ….. ….. ….. pg 33

Practice Multiple Choice ….. ….. ….. ….. ….. ….. ….. ….. pg 34

Chemistry Notes

Praxis Chemistry Glossary of Terms

A Acetones: Ketone hydrocarbons with 2 carbons; standard naming now uses ethanone. Acid Dissociation: The dissolution of an acid into aqueous solution which can also be described as a proton transfer. Acidity: How “acid” a substance is. Acidity for oxides increases down and to the right across the periodic table. (See alkalinity) Acids: less than 7 on the pH scale: (See Lewis Acids, Bronsted-Lowry Acids and Arrhenius acids). Covalent oxides with high oxidation states and high electronegativity are strong acids. Acids have a sour taste, dissolve metals and turn litmus paper red. ****Acid/Base Strength Trends: The strongest acid in a series with the same central atom is the acid with the central atom at the highest oxidation number. The strongest acid in a polyprotic series is the one with the most protons. The strongest acid in a series with different central atoms at the same oxidation number is usually the central atom at the highest electronegativity. Acid-Dissociation Constant: the equilibrium constant for the ionization of a weak acid to a hydrogen ion and its conjugate base (See + - Conjugate Acids/Bases). Ka = [H ][X ] / [HX], where HX is the weak acid. (See Acid/Base Dissociation Constants examples) Activated Complex: An activated complex is a collection of intermediate structures in a chemical reaction. This structure persists while bonds are breaking and new bonds are forming. Activation Energy: Activation energy (Ea) is the minimum energy to overcome the barrier to the formation of products. The formula can be written: ΔE = Eproducts – Ereactants; Ea = Eactivated complex – Ereactants. Alcohol: In the hydroxyl group, a hydrocarbon with a OH-1 ion. Name-ol, example propanol Aldehydes (Carbonyl group): one end carbon bonded to Hydrogen and a double bond to Oxygen. Name- al or name-aldehyde (old); for example, propanal or propionaldehyde. Aldose: containing a aldehyde group Alkali Metals (group 1): 1 valence electron (far from nucleus). These metals form weak metallic bonds and have low melting points; larger radius increases weakness of bonds. Appearance: soft, shiny, low density. Properties tend to decrease as atomic number increases. Alkaline: (See Bases) Alkalinity: How “base” a substance is. Oxides alkalinity increase up and to the left on the periodic table. (See acidity) Alkaline Earth Metals (group 2): 2 valence electrons: stronger than alkali metals. Appearance: grey, metallic solids. Alkanes: hydrocarbons with single bonds between carbons: can be straight chain or ring (cyclo-). Alkenes: hydrocarbons with at least one double bond between carbons and no triple bonds. Alkynes: hydrocarbons with at least one triple bond between carbons. Allotropes: Different forms of the element in the same phase 4 Alpha Particle: ( ) a particle of nuclear decay, equivalent to a helium nucleus ( 2 He) that is released when an unstable nucleus spontaneously releases energy. Alpha radiation has an ionization of +2 and low penetration. Amino: A hydrocarbon molecule with an ammonium group. Amino Group: both amine and amide. Amines C-NH2-H: Amides C double bonded to an Oxygen and single bond to a NH2 ion Anhydride Acid: Carbonyl (oxycarbon) group, similar to an ester but both adjacent carbons are double bonded to oxygen. Name-oic anhydride. Anion: negative ion Annihilation: positron + electron = 2 gamma photons Aromatic Hydrocarbons: Benzene molecules (C6H6) (cyclo-hydrocarbons) with a flat orientation and bonding electrons in the π-orbital are shared by all atoms in the molecule (free moving). Carbon molecules bonded with alternating double and single bonds. (See naming rules for hydrocarbons for additional information.) Arrhenius Definition of Acids and Bases: acids free H+ ions, bases free OH- ions (hydroxide) only. This is the most specific of all acid/base definitions. Atomic Radius: progression from small to large diagonally down to the left. Aufbau Principle: building up rule – filling orbitals at lowest energy. Avogadro’s Law: the number of molecules or atoms in a specific volume of gas is independent of their size or the molar mass of the gas. 23 Avogadro’s Number: The number of molecules per mole usually given as NA = 6.022 x 10 .

Chemistry Notes

B Base-Dissociation Constant: the equilibrium constant for the addition of a proton to a weak base by water to + - form its conjugate acid and a hydroxide ion (See Conjugate Acids/Bases). Kb = [X ][OH ] / [weak base]. (See Acid/Base Dissociation Constants examples) Bases: more than 7 on the pH scale. (See Lewis Bases, Bronsted-Lowry Bases and Arrhenius Bases.) Strongest bases are largest size and lowest metal ion charge. Bases taste bitter, feel slippery and turn litmus paper blue. Benzene: (See Aromatic Hydrocarbons) 0 Beta Particle: ( ) A particle of nuclear decay: Beta decay can be positive (conversion of a proton to a neutron and a positron 1 e), or 0 negative (conversion of a neutron to a proton and an electron -1e). This process also frequently involves the release of a neutrino. Beta has an ionization charge of -1 or +1. Boiling: The phase change process going from liquid to gas; also called vaporization. Boiling Point Elevation: (See Colligative Properties) Boyle’s Law: There is an inversely proportional relationship between pressure and volume in an isothermally closed gas system. Bronsted-Lowry Acid/Base: Acids transfer protons to bases regardless of the nature of the anion accepting the proton (more general than the Arrhenius definition). Bronsted acids and bases exist in conjugate pairs with and without a proton. (See Conjugate Pairs and Bronsted-Lowry acid/base reactions). If two or more acid/base conjugate pairs are present, the stronger acid will transfer a proton to the conjugate base of the weaker acid. Buffer Solutions: A solution that resists a change in pH after the addition of small amounts of an acid or base. Buffer solutions require the presence of an acid to neutralize an added base and the presence of a base to neutralize an added acid. The two components must not neutralize each other; a conjugate acid-base pair is present in buffers. They are prepared by mixing together a weak acid or base and a salt of the acid or base. (See Buffer Solutions in examples). Buffering Capacity: The amount of acid/base that a buffer solution can neutralize before large pH changes begin to occur. Buret: Also spelled burrette. A device that allows the slow, gradual addition of a liquid substance.

C Calibration Curve: Use of positive samples or known data to determine the standard response of a measurement device or experimental set-up. Carbohydrates: Hydrocarbon molecules formed from monosaccharides – simple sugars (CH2O). Monosaccharides contain multiple hydroxyl groups and may be aldose (containing a aldehyde group) or ketose (containing a ketone group). Carboxyl Acid: hydrocarbon acid. Carbon on one end bonded to an OH-1 ion and a double bond to an Oxygen. Name-oic acid. Usually written as –COOH Catalyst Reaction: A and B are reactants, C is the catalyst. (A + C → AC then AC + B → AB + C) Net Reaction: A + B → AB Catalysts: a substance that increases the rate of reaction without being permanently changed during the reaction (involved in intermediate reaction steps) and has no impact on the chemical equilibrium. Catalysts reduce reaction energy. Homogeneous catalysts are in the same physical phase as the reactants. Biological catalysts (enzymes) are usually homogeneous catalysts. Heterogeneous catalysts are in a different physical phase as the reactants, such as a solid surface in which molecules of a liquid or gas can temporarily attach. Catalytic converters are heterogeneous catalysts. Cation: positive ion Charles’s Law: There is a direct relationship between temperature and volume in an isobarically closed gas system. Chemical Equilibrium: (See Equilibrium.) Chemical Properties: heat of combustion, heat of reactivity, pH and pOH, electromotive forces, explosiveness and flammability. Chiral: molecules that lack an internal plane of symmetry. cis-trans Isomers: Stereo-isomers where the elements are in the same order but have a different spatial arrangement (bond angles). cis-trans Isomers only occur when π-bonds prevent rotation, such as in alkenes, alkynes, and cycloalkanes. (See Isomers, Structural Isomers, Functional Group Isomers, Position Isomers, Structural Isomers) Colligative Properties: Properties that depend on the number of solute particles present in a solution. Example: the addition of non-volatile solutes to a liquid solvent causes a reduction in the number of molecules that can evaporate – reducing vapor pressure and driving equilibrium toward the liquid phase and increasing boiling point (decreasing melting point). (See Colligative Properties Table). Common problems: consider the number of particles formed. Glucose does not ionize, so 1 mole glucose forms 1 mole of solvent particles. Sodium Chloride splits into sodium cations and chlorine anions so 1 mole sodium chloride forms 2 moles of solvent particles.

Chemistry Notes

Collisions: Important during reactions (See Reaction Rates, Kinetic Molecular Theory, Rate Laws). Molecular collisions required for reactions. Not all collisions occur with sufficient energy for reaction to occur, not all collisions occur with the required molecular orientation. Colliod: a mixture in which the particles are larger than molecule size but too small to be visible. Combination: (See synthesis) Combustion: A reaction process in which a substance combines with oxygen to form carbon dioxide (carbon monoxide) and water. Concentration: the moles of solute per volume of solution; normally referred to as molarity (M). Sometimes given in Normality or Molality. (See Molality and Normality). Conjugate Pairs: According the the Bronsted-Lowry acid/base definition, acids and bases in a reaction exist in pairs between the forward reaction acid and the reverse reaction acid – also true for the bases. (See Bronsted-Lowry example reactions and Bronsted- Lowry acid/base) Conjugated Molecules: molecules with double bonds on adjacent molecules and aromatic (benzenes) that contain 2 or more π bonds on adjacent atoms. Electrons in these molecules are free to move from one atom to the next on the same molecule (delocalized). Control: Non-treated samples in an experiment necessary for comparison to the treated samples. Negative controls are known to lack the effect under observation; positive controls are known to contain the effect. Covalent Bonds (shared electron pairs) between atoms of relatively close electronegativity Covalent Oxides: Form acidic solutions in water. Oxide combines with water to form an acid. Covalent oxides with high oxidation states and high electronegativity are strong acids. Critical Reaction: The mass of uranium/plutonium in a fission reaction is sufficient that each neutron triggers another fission event. Critical Temperature: the point at which no additional Press will keep substance in its liquid state. Critical Pressure = VP at critical temperature (where surface tension = 0). Crystallization: a process in which dissolved solute collides and reforms into visible particles/

Dalton’s Law of Partial Pressure: Ptotal = ΣPi or PtotalV= (ΣPi)V = (Σni)RT Decomposition: a process in which one reactant separates into two or more products. Degenerate Orbitals: in the absence of a magnetic field, orbitals in the same subshell with different ml have the same energy. Delocalized Electrons: (See Conjugated Molecules) Density: mass/volume. Intermolecular forces contribute to density by bringing the nucleus closer (trends are similar to trends in melting point). Dependent Variable: The variable being observed and measured in an experiment. Deposition: The phase change process going directly from gas to solid. Diatomic Elements: Elements found in diatomic form and not as single atoms: Oxygen, Hydrogen, Fluorine, Nitrogen, Chlorine and Bromine and Iodine. Dipole-Dipole Bonds: between polar molecules where partial + attracted to partial -; also called hydrogen bonds. Dipole-Induced Dipole: the partial charge of a permanent dipole can also induce a dipole in a non-polar molecule that is weaker than a similar ion-induced dipole. Directionality of a Reaction: (See equilibrium and equilibrium constants) Process: Insert the non-equilibrium concentrations into an equilibrium expression to obtain the reaction quotient (Q). If Q < keq, then there are too many reactant molecules and reaction favors products. If Q> keq, then there are too many product molecules and the reaction will favor reactants. If Q = keq, then the reaction is at equilibrium. Double Displacement: A process in which two reactants switch components to form two new products. AX + BY → BX + AY. (May not be a redox reaction). Duet Rule: atoms that ionize such that they have a full 1s orbital: H, He, Li and Be

E Effusion (r): escaping gas into region of lower pressure Electrochemical System: (See voltaic pile or voltaic cell) Electrolytic Cell: Use electricity to force non-spontaneous reactions (opposite of voltaic cells). In an electrolytic cell, the reduction occurs at the cathode (red-cat, gains an electron) and the oxidation occurs at the anode (an-ox, loses an electron). Electrolytic System: Uses inert electrodes (such as the electrolysis of pure water – see examples). 4

Chemistry Notes

Electron Affinity: related to electronegativity – attraction to electrons. Halogens have the highest electron affinity. Filled s or p shells have low electron affinity. Electron Configuration: diagram showing the ground state orbitals of all electrons in an atom. Electronegativity: a relative measure of affinity for electrons. Diagonal progression up and right (weak to strong). As difference in electronegativity increases bonds go from nonpolar covalent (ΔEN < 0.4) to polar covalent to ionic (ΔEN > 1.7). Electroplating: a type of electrolytic system in which the electrodes are not inert and the dissolved metal cations leave the anode and travel to the cathode (item being plated). Elementary Reactions: Single elementary reaction (k = 1) a reaction in which the number of molecules determine rate law (especially unimolecular processes, decomposition, decay). Secondary order elementary reactions (k = 2) 2 include bimolecular processes such as composition A + A → A2 where rate = k[MA] or A + B → AB where rate = k[MA][MB]. Most reactions are multi-step processes involving reaction intermediates where the slowest in the series is the rate-limiting step. Empirical Formula: The lowest whole number chemical formula for a compound. Enantiomers: the mirror image isomers of an organic carbon molecule which cannot be created through rotation. Endergonic: non-spontaneous reactions Endothermic: Process in which heat is absorbed. Endothermic reactions are less likely to occur spontaneously than exothermic reactions. Enthalpy: H = U + Wout or H = U + PV; total energy change by adding or removing heat at a constant pressure. JOULES PER MOLE Enthalpy of Combustion is the heat of reaction when a substance burns in O2 to form completely oxidized products at 25°C and 100 kPa and is always exothermic. Entropy: measure of disorder in a system. Equilibrium: Dynamic equilibrium occurs when two opposing (reversible) processes occur at the same rate (forward rate = reverse rate). At a macro-scale, there is minimal observable change. At a micro-scale, there are two balancing events occurring. Chemical reactions do not “go to completion”; products are generated from reactants to the point when the reaction no longer seems to occur (forward = reverse). Equilibrium can be either homogeneous equilibrium or heterogeneous equilibrium. p q m n Equilibrium Constant: Keq = [MR] [MS] / [MA] [MB] . The equilibrium constant will vary with state conditions. The equilibrium expression for a reaction is the reciprocal of the reverse reaction. For heterogeneous equilibrium, the concentration of a pure liquid or solid in moles per liter (M) cannot change; In the case of a heterogeneous reaction, (such as this, where all are gas except water which is a liquid) the concentrations of pure liquids or solids (as opposed to aqueous) are not included. (See example problems for equilibrium constants) Equilibrium Reaction: See also La Chatelier’s Principle. A reaction at equilibrium contains a constant ratio of chemical species (determined by an equilibrium constant). mA + nB ↔ pR + qS. Equivalence Point: The volume at which a titration curve steeply changes; for strong acid/base titrations this occurs at a pH = 7. Ester: hydrocarbon with an oxygen bonded between carbons (single bond) AND a double bonded oxygen to one of the adjacent carbons. Named by (short chain)yl (long chain)-oate Ethers: hydrocarbons with an oxygen bonded between carbons (single bonds). Named by (short chain)yl (long chain)yl ether. Exergonic: Spontaneous reaction Exothermic: Process in which heat is released. Evaporation: Sometimes used interchangeably with vaporization but technical the phase change process in which individual high energy liquid molecules escape without the addition of external energy.

F Families: Also called groups. Families are columns of elements on the periodic table. Fissile Material: Material used in fission reactions. Freezing: The phase change process from liquid to solid; sometimes called solidification. Functional Groups: Name given to the different categories of attachments to organic molecules. (See hydrocarbons) Functional Group Isomers: Molecules with the same empirical formula but arranged in different functional groups.

Chemistry Notes

G Gamma Particle: A gamma particle is a high-energy photon and is both part of ionization nuclear decay radiation and the electromagnetic spectrum. Gamma photons are released from the nucleus of an unstable atom. The gamma particle is the most penetrating of the three types of nuclear decay. Gay-Lussac’s Law: There is a direct relationship between pressure and temperature in an isovolumetrically closed gas system. Gibbs Free Energy: ΔG = ΔH - TΔS Graham’s Law: rate of effusion (r) is inversely proportional to the square root of the molecular mass.

Haber Process: Process that allowed the artificial production of ammonia. N2(g) + 3H2(g) ↔ 2NH3 (g) (Fe catalyst) Halide: hydrocarbon with one of the group 17 on one end. Example: bromobutane. Halogens: Group 17 elements. Often have strong odor. London forces increase as atomic number increases. Heat of Formation: The standard enthalpy of formation or standard heat of formation of a compound is the change of enthalpy that accompanies the formation of 1 mole of a substance in its standard state from its constituent elements in their standard states (the most stable form of the element at 1 bar of pressure and the specified temperature, usually 298.15 K or 25 degrees Celsius). An element in its most stable form has a heat of formation of 0. Heat of Reaction = Heat of Products – Heat of Reactants. If Hproducts > Hreactants heat is absorbed and process is endothermic. If Hproducts < Hreactants heat is released and process is exothermic. Heat of Solution: The energy change associated with the process in which a solute dissolves. 1) Heat to break up the solute (crystal lattice energy) must be absorbed. 2) Heat to form solute-solvent (heat of hydration) is released. 3) The heat of solution is the heat of formation minus the heat of dissolution. Henry’s Law: the solubility of a gas in a liquid at a particular temperature is proportional to the pressure of that gas above the liquid. Hess’s Law: ΔHnet reaction = Hreaction 1 + Hreaction 2 + …… The total enthalpy change will be the sum of the changes for each step. Heterogeneous Catalyst: (see Catalyst) Homogeneous Catalyst: (see Catalyst) Hund’s Rule: before any 2e occupy the same orbital, other orbitals in that subshell must contain one electron. Hybridization: a bonding process in which there is pre-bond promotion of one or more electrons from a lower subshell to a higher energy subshell followed by a combination of the orbital into degenerate hybrid orbitals: occurs for ns2 ns2p1 or ns2p2; also when oxidations #’s are +2, +3, and +4 regardless of family. Hybridization with Incomplete Bonding: Occurs when atoms are bonded by multiple covalent bonds. See Hybridization Rules. Hydrates: water molecules occupying positions within the lattice of an ionic crystal (water of hydration). Hydration: (See Solvation) Hydrocarbon Derivative: the substitution of one or more atoms with unpaired electrons into a hydrocarbon backbone. Hydrocarbon Groups: Alcohols (hydroxyl group); Ketones (Carbonyl group); Aldehydes (Carbonyl group), Carboxyl Acid (hydrocarbon acid). Amino group: both amine and amide. Sulfhydrl group: nonpolar (thio means sulfur and hydrogen). Other groups include ethers, esters, anhydride acids, halides, and nitriles. (See Diagrams (Biologically Important Hydrocarbons) and Naming Hydrocarbons for additional information) Hydrocarbon Reactions: (see Organic Reactions) Hydrogen Bonds: strong dipole-dipole interactions that form between the hydrogen atom of one molecule and a FLUORINE, OXYGEN or NITROGEN of an adjacent atom. (Typically found as intermolecular bonds although some larger molecules may have hydrogen bonds as intramolecular bonds). + Hydronium Ion: H3O or protonated water Hydrophilic: Water loving Hydrophobic: Water hating Hydroxide: a polyatomic ion of OH with an oxidation number of -1. Hypothesis: a proposal of a possible answer to a question; testable. Hypotheses are sometimes called predictions.

I Ideal Gas: a gas that follows the rules of kinetic molecular theory. Primarily the gas is chemically non-reactive. Ideal Gas Law: the relationship between pressure, volume, moles and temperature of an ideal gas. R = 8.31 Pam³/mol K Immiscible: When two liquids do not mix. Independent Variable: The variable being purposefully manipulated in an experiment or study.

Chemistry Notes

Induced Dipole: the distortion in a non-polar molecule caused by the presence of an ion. These are the only forces between nonpolar molecules in solids at low temperatures. These are the weakest type of intermolecular bond and are stronger in large molecules or with larger surface areas since a large cloud is more easily polarized. Induced Dipole Bond: Intermolecular attractions due to induced dipoles in a non-polar molecular; also called London Forces or Van Der Waal. Intermolecular Forces: Dipole-Dipole between polar molecules where partial + attracted to partial -; also called hydrogen bonds. Induced dipole: the distortion in a non-polar molecule caused by the presence of an ion. Intermolecular attractions due to induced dipoles in a non-polar molecular; also called London Forces or Van Der Waal. Intramolecular Bonds: Covalent (shared electron pairs) between atoms of relatively close electronegativity, Ionic (electrostatic attraction between ions), metallic (electrons commonly shared). Ion-induced Dipole: when a non-polar molecule or noble gas atom encounters an ion, the electron density is distorted and creates an induced dipole. Ion-dipole Interactions: In solutions, a dissolved salt (ions) will be separated by the partial charge of the polar covalent molecule (such as water). Ion-Product Constant: Kw is an equilibrium constant relating to the acidity/alkalinity of pure water (see pH, pOH and examples on + - -14 acidity). –log Kw = -log[H ] – log[OH ] = -log(10 ) so therefore pH + pOH = 14. Ionic Bonds (electrostatic attraction between ions). Ionic bonds form brittle compounds with high melting points that are not conductive in solid form (are conductive in solution). Formation of ionic compounds is EXOTHERMIC. Ionic compounds are generally stable but will react if the reaction forms a more stable salt. Ionic Characteristics: Increase up and to the left on the periodic table for halide compounds (compounds with Group 17 elements) Ionic Equation: balanced equation written in terms of the positive (cations) and negative (anions) ions. Ionization Energy: Increases across the periodic table up and to the right. Increases with greater nuclear charge and as sub-shells fill. Decreases with shielding effect (shells between nucleus and outer electrons) and with atomic radius. Isomers: molecules with the same empirical formula but different structure. (see cis-trans Isomers, Structural isomers, Chain Isomers, Functional Group Isomers, Position Isomers) Isobaric: A system in which there is no change in pressure. Isothermal: A system in which there is no change in temperature. Isovolumetric: A system in which there is no change in volume. IUPAC : international union of pure and applied chemistry

K Ketones: Carbonyl group hydrocarbons have a double bonded oxygen (usually to one of the center Carbons). Name-one, example proponone (old name: acetone). Ketose: containing a ketone group Kinetic Energy: The energy of motion: the ability of a moving object to exert a force/do work/or exchange energy with its surroundings. (KE = ½mv2) Kinetic Molecular Theory: 1) large number of molecules moving independently and randomly with no intermolecular bonds; 2) KE = U (no PE) and KEavg is directly proportional to Temperature. 3) Energy can be transferred during collisions with no lost KE. 4) Molecule size is negligible compared to space occupied. Related to reaction rates (see Reaction Rate).

L La Chatelier’s Principle: If a reaction at equilibrium is disturbed, changes in the system will occur to reestablish equilibrium in a manner that counteracts the disturbance. Reactions at equilibrium contain a constant ratio of chemical species. Adding a substance will shift the reaction so a new equilibrium is established by consuming some of the added substance. Increasing pressure in a system will push the process in a way to reduce the moles of gas. When temperature changes for an isobaric process, the reaction shifts in the direction that absorbs heat (pushes the endothermic process). Law of Mass Action: The mathematical relationship between the concentration of reactants and products in a system. Lewis Definition of Acid/Base: Bases donate unshared electron pairs to acid, regardless of whether the donation is made to a proton or another atom. (See Lewis Acid/Base in diagrams and examples) Lewis Dot Structure: model showing valence electrons as dots to indicate covalent bonds.

Chemistry Notes

Limiting Reagents: the reactant present in lowest value (compared to balanced equation) that determines the possible amount of product generation. Lipids: (Fatty Acids) Large hydrocarbon molecules which are insoluble in water. Lipids can be complex (triacylglycerides) which contain three ester groups or simple (cholesterol) which have no ester groups. Saturated fats have single bonds chain carbons; unsaturated fats have one or more alkene group. Lipids have a non-polar side (hydrophobic) and a polar side (hydrophilic). London Forces: also called induced dipole or Van Der Waal London Dispersion Forces: an interaction between an induced-dipole and an induced-dipole. A non-polar molecule may temporarily induce a weak dipole in an identical non-polar neighbor due to the instantaneous position of its electrons. The London dispersion force is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles. This force is sometimes called an induced dipole-induced dipole attraction. London forces are the attractive forces that cause nonpolar substances to condense to liquids and to freeze into solids when the temperature is lowered sufficiently.

M Maxwell-Boltzmann Distributions: Show the effect of temperature on velocity distribution. Melting: The phase change process from solid to liquid. Also called liquefaction. Metal Oxides: metallic atoms bonded to oxygen atoms. These compounds typically form basic solutions in water. Metallic Bonds (electrons commonly shared). Also termed “electron sea”. Metallic bonds form good conductors, lustrous, malleable and ductile. Metalloids: Smallest category of the periodic table; usually include Boron, Silicon, Germanium, Antimony, Tellurium, Astatine. Miscible: When two liquids are capable of mixing. Mixture: A combination of two of more substances that are not combined chemically – physical change only. Model: A representation of a system, theory, etc used to help illustrate its behavior. Molality: A measure of concentration given in moles solute per mass (in kg) of solvent. (See concentration and Molarity). Molality does not change with temperature (due to volumetric expansion). Molar Heat Capacity: the heat required to change the temperature of 1 mole by 1°C. Molarity: (See concentration). Mole Fraction: number of moles of a given component divided by total number of moles. (Xa) Molecular Mass: mass of a molecule; can be found by ρV/n where n is number of moles. Molecular Orbital Theory: Shows approximate electron density shapes by combining atomic orbitals. This is primarily used in discussion of sigma and pi orbitals.

N Network Solid: a solid made up of covalent bonds that act as a large single molecule. These are hard, strong and have a high melting point. Neutralization: Also called acid/base neutralization. A reaction process in which an acid and base are combined to form salts. The reaction equation looks different for Arrhenius, Bronsted-Lowry and Lewis acid/bases due to the difference in what is defined as an acid and a base. Nitrile: hydrocarbon with a cyanide ion at one end (final carbon triple bonded to nitrogen). Name-onitrile. Noble Gases: Group 18 atoms. Usually inert (non-reactive) but some of the heavier noble gases may bond with fluorine or oxygen. Noble Metals: Also called coinage metals. Group IB metals Non-polar: A bond in which the balance of partial charge is 0. Example: Methane (CH4) is non-polar. Normality: A measure of concentration given in “equivalents” solute per volume of solution. An equivalent is a whole number multiplier of moles solute defined so that 1 equivalent of one reagent reacts with 1 equivalent of the other reagent. Nuclear Isomerism: the same nucleons may arrange themselves within a nucleus in different ways with different energies. (ie release of gamma rays) Nucleic Acids: Nucleic acids are formed from nucleotides which are made up of a phosphate, a sugar and an amine base. These molecules can be linked together to form ribonucleic acid (RNA) and deoxyribonucleic acid (DNA). Nucleophile: A reagent that forms a chemical bond to its reaction partner (the electrophile) by donating both bonding electrons.[

O Octet Rule: Atoms that ionize such that they have a full s-p orbital of valence electrons.

Chemistry Notes

Orbitals: s (2), p (6), d (10), f (14) and g (18): 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s…….. Organic Molecules: Hydrocarbons Organic Reactions: Include addition, substitution, and elimination. Addition: Two hydrocarbons will react when the double bond of an alkene is replaced with single bonds (replacing a -bond with a -bond). Substitution: The replacement of a group in a hydrocarbon, including SN1 or SN2 type reaction. Elimination: when a functional group is eliminated from the adjacent carbons; this reaction is favored by the presence of a strong base. (See diagrams for Organic Reactions for details) Osmosis: a flow of liquid through a semi-permeable membrane separating a dilute and concentrated solution. The flow is from concentrated to dilute until equilibrium is reached. Osmotic Pressure: The pressure required to prevent the osmosis from a pure solvent (no solute) into a solution. It is proportional to molarity of the solution and is a colligative property. (See colligative property). Oxidation Number: The charge on an ion. Oxidation numbers typically the number of electrons that must be lost/gained to achieve full outer shell. Oxidation numbers assigned with the assumption of ionic bonds. Some elements can have various oxidation numbers. Oxidation Rules: Free elements (including diatomic) are assigned 0. Free monatomic ions are assigned an oxidation number equal to the charge of the ion. Group 1 elements in compounds are assigned +1; group 2 elements in compounds are assigned +2. Zinc (Zn) and Cadmium (Cd) are assigned +2 in compounds; Aluminum (Al) and Gallium (Ga) are assigned +3 in compounds. Hydrogen is assigned +1 in compounds except in hydrides in which it is assigned -1. Oxygen is assigned -2 in compounds except in peroxides where it is assigned -1. Fluorine is always assigned -1 in compounds. Oxidize: A process in which a substance (element or polyatomic ion) increases its oxidation number. Oxidizing Agent: a substance with the ability to oxidize other substances by removing electrons. The oxidizing agent is reduced.

P Pauli Exclusion Principle: no 2 electrons have the same quantum numbers Percent Composition: Percentage of a given substance to the total – usually by mass. Percent Yield: Actual yield divided by theoretical yield for a reaction. Periods: Rows of elements on the periodic table. pH: the negative base-10 logarithm of the hydrogen-ion molarity (concentration). A measure of the relative acidity (pH<7) or alkalinity (pH>7). (See also pOH) pH = -log[H+]. An increase of concentration by a factor of 10 leads to an increase in pH of 1. Photon: the term given to the particle nature of light, consists of a small “packet” of energy. Physical Properties: color, mass, smell, melting point, boiling point, density. Usually result from strength of the intermolecular forces. Pi Orbitals (π): Formed by bonds which lie in regions separate from a line drawn between two atoms in a bond. Found in double and triple bonds. Pi orbitals prevent rotation along the bond axis. pOH: the negative base-10 logarithm of the hydroxide-ion molarity (concentration). pOH = -log[OH-] Polar: a molecule in which the charge distribution is not uniform. To determine polarity: diagram the molecule, assign EN values to atoms, and determine cancellation of any polarities. Assign a net dipole. Polyprotic Acids: Acids that can transfer more than one proton (See Bronsted-Lowry acid/base) Polyprotic Series: a reaction chain for weak Bronsted-Lowry acid/base pairs. Position Isomers: Molecules with the same empirical formula that differ by the attachment position of elements within their functional group. Precipitation Reaction: a reaction type in which aqueous solutions combine to form a solid precipitant which falls out of solution. Protonated Water: term for H+(aq) Proteins: Hydrocarbon molecules formed from amino acids which contain both an amine group and a carboxylic acid. Proteins can be dipeptide (two amino acids) or polypeptide (protein) with more than two amino acids. Proteins include keratin, insulin, and enzymes.

Q Qualitative Data: Descriptive, non-numerical, observations Quantitative Data: Data collected through observational measures and compared to a scale. 2 Quantum Numbers: n, l, ml describe orbitals; n is the principle quantum number related to the distance from the nucleus 2n = shell (orbital); l is azimuthal quantum number between 0 and n-1 describes the angular momentum of the system. ml is the magnetic

Chemistry Notes quantum number between –l and l. With a single electron, n is the only relevant number; in a multi-electron system the inner electrons shield electrons further away and reduce the forces acting on them.

R Raoult’s Law: The vapor pressure of a solution with non-volatile solutes is equal to the mole fraction of the solvent multiplied by the v v vapor pressure of the pure solvent. P solution = P pure solvent (mole fraction)solvent. (See Colligative Properties) Rate Laws: an equation relating reaction rate to concentration. Rate laws cannot be predicted by Stoichiometry. Rate Law Constant: the constant (k) used in determining reaction rates and varies with overall reaction order. k = 0, where k is in molarity per second (M/s) for reactions such as decay (or the negative slope of a reactant/time graph). k = 1, where k is in s-1, for reactions directly proportional to concentration (or the negative slope of ln(reactant)/time graph). k = 2, where k is in (M•s)-1 for reactions proportional to the square of the concentration (or the negative slope of (reactant)-1/time. Rate Law Formula: rate = k[reactant 1]a[reactant 2]b, k is typically used as the rate constant. a, b are reaction orders where 0 means the concentration has no effect as long as the reactant is present, 1 means that reaction rate is directly proportional to concentration, and 2 means that the rate is related to the square of the concentration. Reaction Mechanism: series of elementary reactions that cannot be determined from Stoichiometry. Reaction Rate: given in molarity per second or found by change in concentration per change in time (ΔM/Δt) or change in product/reactant per time. Derivative of product change per time; assumption is linearity. Reaction rate constants are directly related to temperature. (See Kinetic Molecular Theory: Increase Temperature – Increased Collisions – Collisions occur with more energy). Reaction Types: Reactions fall into the following categories: Combination (Synthesis), Decomposition, Single Substitution (Single Displacement or Single Replacement), Double Substitution (Double Displacement, Double Replacement, Ion Exchange or Metathesis) and Isomerism. Reactivity: The reactivity of substances is based on stability considerations that are often localized to the most polarized regions. Examples: exposed surfaces of alkali or alkaline earth metals rapidly oxidize; halides easily form oxides and combine with other halogens. Strongest reactivity is between substances with very different electronegativities. Reactivity Series: Metal (reaction with O2) K•Na(burn violently) •Ca•Mg•Al•Zn(Burns rapidly) •Fe•Pb•Cu•Hg(oxidize slowly). Redox Reactions: A redox reaction is a reduction/oxidation reaction in which substances increase (oxidize) or decrease (reduce) their oxidization numbers. Reducing Agent: a substance with the ability to reduce other substances by adding electrons. The reducing agent is oxidized. Reduction: A process in which a substances (element or polyatomic ion) drop to a lower oxidation number. Replacement: See Single Replacement and Double Replacement. Resonance: Equivalent Lewis structure resonance forms, such as O3-ozone (one double bond and one single bond – bonds actually exist in resonance between the two states).

S Saturated: The condition in which the dissolved and undissolved solute particles in a solution are at equilibrium and no further solute can be dissolved. Sigma Bonds (σ): occurs with a bond between 2 s subshell orbital electrons. At least 1 electron pair in every bond is a sigma bond. Sigma bonds can also be formed by hybrid sp and p atomic orbits IF the orbits overlap such that the axis between the bonded atoms runs through the center of the combined electron density. Single Replacement: A process in which the part of one reactant is replaced by the other reactant. A + XB → B + XA. Occurs when there is a greater electronegativity attraction between A and X than between B and X. (Redox reaction) Solubility: “like dissolves like”. Nonpolar materials have intermolecular bonds with weak London Forces that cannot overcome the hydrogen bonds in water (polar solvent). Nonpolar solutes dissolve more readily in nonpolar solvents; polar solutes dissolve more readily in polar solvents. Electrolytic materials dissolve and ionize; nonelectrolytic materials dissolve without ionizing. (See solubility rules in Tables). Factors affecting solubility: Gas: ↑partial pressure of a gas will ↑solubility of that gas in a liquid; temperature will ↓solubility of a gas. Liquids and Solids: endothermic processes: ↑temperature will ↑solubility, exothermic processes: ↑temperature will ↓solubility. (Aqueous salts are usually endothermic). Solubility Product Constant: The equilibrium constant for an ionic solid in contact with a saturated aqueous solution. Ksp = [cation concentration]p[anion concentration]q (p = moles cation, q = moles anion). Solute: Substance dissolved in a solvent. Solution: The physical mixture of two substances in which no chemical change is taking place. Solutions form when the intermolecular attractive forces between the solute and the solvent are about as strong as those in the solute or solvent alone. 10

Chemistry Notes

Example: NaCl and H2O: Water molecules interact with the sodium cations and the chlorine anions to overcome the attraction between them in their crystal form. Na+ and Cl- interact with the water molecules to overcome their attraction to each other. Solvation: the intermolecular attraction between solute and solvent molecules. If the solvent is water, this is also called hydration. Solvent: a liquid, solid, or gas that is capable of dissolving another liquid, solid or gas. Specific Heat: same as specific heat capacity. Q = mCΔT Spectator Ions: ions present but not involved in changes of oxidation number in a redox reaction. Spontaneous Reactions: –ΔH (exothermic) and +ΔS cause a decrease in energy and an increase in entropy o Standard Cell Potential: (E cell) The voltage generated by an electrochemical cell at 100 kPa and 25°C, when all components of the reaction are pure materials or solutes at 1 M. The standard cell potential is calculated from the sum of the two half-reaction o o o + potentials: E cell = E reduction/cathode - E oxidation/anode. All half reaction potentials are relative to the reduction of H to from H2. The cathode is the terminal with the more positive voltage and the anode is the terminal with the more negative voltage. Standard Molar Volume: 1 mole of gas at STP occupies 22.4 L Standard Thermodynamic Value: occurs at 25°C and 100 kPa Stoichiometry: a branch of chemistry that deals with the relative quantities of reactants and products in chemical reactions. Reactions can be examined mole:mole, mole:mass, and mass:volume. STP: standard temperature and pressure, 0°C and 1 atm Strong Acid/Base: Strong acids and bases completely dissociate in water; they are also strong electrolytes. The statement “strong” acid gives no indication of the potential health hazard associated with the acid/base. Strong acids pass protons to water making hydronium more readily than hydronium transfers protons back. (See acid/base table). Structural Isomers: Isomers of chain molecules with the same empirical formula that differ primarily by their branching. Sub-Critical Reaction: A fission reaction in which too many neutrons are lost and the chain reaction stops. Sublimation: The phase change process that goes directly from solid to gas. Surface Tension: balance of weak attraction of molecules to vapor and strong attraction of molecules to liquid. Surface Tension inversely related to temperature. Supercritical Fluid: temperature and pressure greater than critical. Fluid has density and bonds like a liquid but expands like a gas. Supercritical Reaction: A fission reaction with sufficient mass of fissile material causes the explosive release of energy. Synthesis: Reactions in which two reactants combine to form one product. Also called combination.

T Theory: An explanation for a body of facts or evidence that can be used to predict future outcomes. Thio: means sulfur and hydrogen such as mercaptan. Titrant: The reagent of known concentration in a titration. Titration: A process of monitoring an acid/base neutralization reaction by use of a pH meter or pH indicator. Titration usually involves determining the molarity of an unknown acid or base through the addition of a acid/base of known molarity. Calculations of molarity assume that one mole of both solutions yields the same number of equivalents. (See Titration in examples). Titration Curve: a plot of a solution’s pH charted against the volume of added acid or base. The shape of a titration curve will indicate the strength of the acid/base and the presence of polyprotic acids. (See Titration Curve diagrams, polyprotic acid) The equivalence point of a titration curve of week acids or bases will not occur at a neutral pH (7). (See equivalence point). Triple Point: the point on a P-T diagram (showing phases of matter) at which the solid, liquid, and gas regions meet. See also critical point.

U Unit Cell: A small box containing one or more atoms, a spatial arrangement of atoms. Unsaturated: The condition in a solution in which the solvent is still capable of dissolving further amounts of solute.

V Valence Shell Electrons: electrons in outermost shell (involved in intramolecular bonds). The most stable atoms have a full valence shell (Group 18 and (lesser) Group 2 (filled s subshell). 11

Chemistry Notes

Valence Number: number of electrons in outermost shell; electrons available for bonding. Van Der Waal: also called induced dipole or London Forces Van ‘t Hoff Factor (i): Unit-less constant directly associated with the degree of disassociation of the solute in the solvent; a factor predicting the change in boiling point or freezing point of a solution after a solute has been added. Substances that do not ionize: I = 1 (sugars); substances that ionize into 2 ions: i = 2 (NaCl); substances that ionize into 3 ions: i = 3. Vapor Pressure is directly related to Temperature until VP = Pexternal (at BP); BP then increases as Pexternal increases. In a closed system with liquid incompletely filling area, molecules at the surface with greatest KE will evaporate, increasing pressure due to gas in system until equilibrium is reached. Partial pressure of substance in a system is at saturated vapor pressure. Vaporization: phase change between liquid and gas. Increasing pressure favors liquid state. Vaporization is not the same as evaporation. Viscosity is ability to flow. Higher Viscosity means higher intermolecular forces; viscosity inversely related to temperature; most liquids, viscosity is independent of pressure (except at very high pressure). Voltaic Cell: A electrochemical reaction that occurs spontaneous and generates potential difference. The reduction occurs at the cathode which acts as the positive terminal and the oxidation occurs at the anode which acts as the negative terminal. Voltaic Pile: (See voltaic cell) VSPER Model: model describing predicted molecular geometry (Valence Shell Paired Electron Repulsion). Electron pairs around the central atom of a molecule repel each other. When some pairs are unshared, the bond geometry is distorted away from unshared pairs because the unshared pairs are pulled in towards the central nucleus. Process: 1) Lewis Dot Structures and 2) Determine number of unshared pairs around central atom. See VSPER Model diagrams.

W Weak Acid/Base: A weak acid or base do not completely dissociate in water; these substances are also weak electrolytes. Weight Percentage: The percentage mass of a component per mass of the solution.

Chemistry Notes

Equations/Variables Used in Common Calculations

Diffusion: 0.5 drate α (T/m) Dilution Equations: MiVi = MfVf Entropy: ΔS = ΔQ/T (Joules per Kelvin) if reaction is defined by the equation xA = yB = x’P + y’Q; then ΔS = x’SP + y’SQ – xSA - ySB Enthalpy: H = U + Wout or H = U + PV Equilibrium Constant: p q m n Keq = [MR] [MS] / [MA] [MB] Gibbs Free Energy: ΔG = ΔH - TΔS Half-Life Equation: n N = N0(½) , where n = t/t0.5 and t0.5 = half life period Heat Equations: ΔHreaction = Hproducts - Hreactants ΔHfusion = enthalpy of fusion (Joules per mole) ΔHvaporization = enthalpy of vaporization (Joules per mole) ΔHcombustion = enthalpy of combustion is the heat of reaction when a substance burns in O2 to form completely oxidized products at 25°C and 100 kPa and is always exothermic. (Joules per mole) ΔHformation = required when elements react to form a compound (also called enthalpy of formation). In most stable form, elements are assigned a value of 0 J/mole Q= mc T (heat change of a substance within a phase, where C = specific heat)

Q= mHv or Q= mHf (heat change of a substance during a phase change where Hv is heat of vaporization and Hf is heat of fusion) Frequently heat equations are in terms of moles instead of mass. Ideal Gas Law: PV = nRT or PV = NkT Partial Pressure: Ptotal = ΣPi pH and pOH pH = -log[H+] pOH = -log[OH-] Raoult’s Law: (See also Colligative Properties Table) v v P solution = P solvent (mole fraction)solvent v v ΔP = - P solvent (mole fraction)solute Rate Law Formula: rate = k[reactant 1]a[reactant 2]b Solubility Product Constant p q Ksp = [cation concentration] [anion concentration] Standard Cell Potential E°cell = E°cathode + E°anode Titration: [A](VA) = [B](VB) Weight Percentage: (mass component / mass solution) x 100

Equation Constants Avogadro’s Number N = 6.022 x 1023 particles per mole Ideal Gas Law Constant R = 8.31 Pa m3 mol-1 K-1 = 0.08206 atm L mol-1 K-1

Chemistry Notes

Example Problems

Acid/Base Dissociation Examples: The base-dissociation for ammonia at 25°C is 1.8 x 10-5. What is the concentration of hydroxide ions in an ammonia solution at equilibrium containing 0.2M ammonia at 25°C? Write dissociation equation + - + NH3 (aq) + H2O (l) ↔ NH4 (aq) + OH (aq) (the weak base (NH3) accepts a proton and forms the conjugate acid NH4 ) Set-up the dissociation constant equation and solve; the concentration of the ammonium ions and hydroxide ions is equal (x). - Kb = [conjugate acid][OH ] / [weak base] 1.8E-5 = [x][x] / 0.2M x² = 3.6 x 10-6 x = 1.9 x 10-3M for OH-

Acid/Base Ion Product Constant What is the concentration of hydroxide ions in an aqueous solution with a hydrogen ion concentration of 2.5 x 10-6M? What is the concentration of hydrogen ions when pure water reaches equilibrium? Write the Ion Product Constant equation where the ion product constant is 1 x 10-14. -14 + - Kw = 1.0 x 10 = [H ][OH ] [OH-] = 1.0 x 10-14 / [H+] = (1.0 x 10-14) / [2.5 x 10-6M] = 4.0 x 10-9M At equilibrium [H+] = [OH-] -14 + - Kw = 1.0 x 10 = [H ][OH ] = x² x = 1.0 x 10-7M

Acid/Base Neutralization for Arrhenius, Bronsted-Lowry and Lewis definitions. Arrhenius: + - 0 - - General: HNO3 + NaOH → NaNO3 + H2O Reduction: H + e → H Oxidation: OH → OH + e The hydrogen cation (+1) combines with an electron to form hydrogen with oxidation number of zero. The hydroxide anion (-1) releases an electron to form a hydroxide with an oxidation number of zero.

Dissociation of an acid: HCl (aq) → H+(aq) + Cl-(aq) + - Proton Transfer form of same reaction: HCl (aq) + H2O → H3O (aq) + Cl (aq)

Bronsted-Lowry: + - + - General: HNO3 + KCN → HCN + KNO3 Ionic Equation: H + NO3 + K + CN → The nitric acid combines with the potassium cyanide salt to form hydrogen cyanide (prussic acid) and potassium nitrate. Not a redox reaction since the oxidation states of the ions do not change. The hydrogen loses its bond to anion on the acid bonds to the anion on the base. The cation from the salt then bonds to the anion from the acid.

If two or more acid/base conjugate pairs are present, the stronger acid will transfer a proton to the conjugate base of the weaker acid.

- + HF (aq) + H2O (l) ↔F (aq) + H3O (aq) The hydrofluoric acid transfers a proton to the water forming a fluoride ion and a hydronium ion. By the Bronsted-Lowry definition, the hydrofluoric acid has a conjugate acid in the hydronium, the water acts as a base with a conjugate base of the fluoride ion.

+ - NH3 (aq) + H2O (l) ↔ NH4 (aq) + OH (aq) The water (acid) transfers a proton to the ammonia (base) forming an ammonium ion (conjugate acid) and hydroxide ion + (conjugate base). NH3 : NH4 pair

- - H3PO4 (aq) + HS (l) ↔ H3PO4 (aq) + H3S (aq) The phosphoric acid (acid) and the hydrogen sulfide ion (base) form a hydrogen phosphate ion (conjugate base) and - - hydrosulfuric acid (conjugate acid). H3PO4: H3PO4 conjugate acid-base pair and HS :H3S is the other.

Chemistry Notes

Lewis: General: An acid joins to a lone electron pair of a base and forms a Lewis Complex

Acid/Base pH and pOH Examples: Examples for calculation of acidity or alkalinity for solutions + - -14 –log Kw = -log[H ] – log[OH ] = -log(10 ) pH = -log[H+] pOH = -log[OH-]

An aqueous solution has an H+ ion concentration of 4.0 x 10-9. Is the solution acidity or basic? Calculate the pH and pOH. pH = -log[4.0E-9] = 8.4, so basic pOH = 14 – 8.4 = 5.6

Acid/Base Strength Examples: The strongest acid in a series with the same central atom is the acid with the central atom at the highest oxidation number. The strongest acid in a polyprotic series is the one with the most protons. The strongest acid in a series with different central atoms at the same oxidation number is usually the central atom at the highest electronegativity. Arrange the following in weakest to strongest: H2SO3, H2SeO3, H2TeO3 → electronegativity of S>Se>Te so: weak H2TeO3, H2SeO3, H2SO3 strong HBrO, HBrO2, HBrO3, HBrO4→ oxidation number highest for HBrO4 so: weak HBrO, HBrO2, HBrO3, HBrO4 strong HI, HBr, HCl, HF → electronegativity of F>Cl>Br>I but HF is a WEAK acid - 2- 2- - H2PO4, H2PO4 , H2PO4 → H2PO4 has the most protons so: weak H2PO4 , H2PO4 , H2PO4 strong

Buffer Solution Description A buffer solution is prepared by mixing together acetic acid (HC2H3O2) and sodium acetate (NaC2H3O2) which is an acetate ion - + + - (C2H3O2 ) and a sodium ion as a spectator (Na ). The equilibrium reaction is: HC2H3O2 ↔ H + C2H3O2 . If an acid is added to the solution, and more hydrogen ions are added to the solution, the equilibrium is forced to the left. If a base is added to the solution, so that hydrogen ions are consumed, the equilibrium is forced to the right.

Equilibrium Constant: Homogeneous: 2 2HI (g) ↔ H2 (g) + I2 (g). keq = [MH2][MI2]/[MHI] Hydroiodic Acid disassociates into hydrogen and iodine gas Heterogeneous: Full equation: Cu(s) + 2Ag+(aq) ↔ Cu2+(aq) + 2Ag(s). Half equations Cu ↔ Cu2+ + 2e- and 2[Ag+ + e- ↔ Ag] 2 keq = [MCu2+]/[MAg+] Copper and an aqueous solution of silver cations react to form an aqueous solution of copper cations and silver.

Gas Laws At STP, 0.250 L of an unknown gas has a mass of 0.429 g. Is the gas SO2, NO2, C3H8 or Ar? Determine the moles of the gas present using ideal gas law, where STP means P = 1 atm (101,300 Pa) and T = 273 K PV = nRT so n = PV / RT = (1 atm)(0.250 L) / (0.08206 atm L mol-1 K-1)(273 K) = 0.011 moles Determine the molar mass of the gas Molar mass = grams / mole = 0.429 g / 0.011 moles = 38.4 g/mol Compare to the molar mass of the given substances: SO2 (64 g/mol), NO2 (46 g/mol), C3H8 (44 g/mol) or Ar (40 g/mol) Most likely Argon.

Heat Flow and Entropy: What is the change in energy and entropy of 10.0 g gold at 25˚C when it is heated to beyond its melting point to 1,300 ˚C using the following data: Solid Heat Capacity of Au = 28 J mol-1 K-1, Molten (Liquid) Heat Capacity of Au = 20 J mol-1 K-1, Enthalpy of Fusion = 12.6 kJ mol-1, and the melting point of Au = 1064ºC Determine the number of moles of gold Molar Mass of Au = 197 g/mol, so 10.0 g / 197 g/mol = 0.051 moles

Chemistry Notes

Determine the energy flow (in Joules) required to heat the solid to the melting point, melt the solid and then heat the liquid to the final temperature. Since a T of 1 K = 1˚C, it is not necessary to covert temperatures at this point. Q = Qsolid + Qmelt + Qliquid = Q = (0.051mol)(28J mol-1 K-1)(1064 – 25˚C) + (0.051mol)(12600J mol-1) + (0.051mol)(20J mol-1 K-1)(1300 – 1064 ˚C) Q = 2367.012 J or 2400 J

Determine the Entropy change of the system where S = Q/Tavg S = 2400 J /((0.5)(1300+25)) = 2400 J / 662.5 K = 3.6 J/K

Heat of Formation: Find the standard heat of formation (ΔHf°) for ethylene where 2C (graphite) + 2H2(g) → C2H4(g); Determine if C(graphite) and H2(g) will react to form ethylene at 25°C and 100kPa; Find the ΔSf° for formation of C2H4(g). Given: ΔHf° for CO2 = -393.5 kJ/mol•C°, ΔHf° for H2O = -285.9 kJ/mol•C°, and ΔHc° for C2H4 = 1411.2 kJ/mol•C° Given: S° for C(graphite) = 5.7 J/mol•K, S° for H2(g) = 130.6 J/mol•K, and S° for C2H4 (g) = 219.4 J/mol•K Write Combustion and Formation Equations in terms of 1 mole formed or combusted: Equation 1: C(graphite) + O2(g) → CO2(g) Equation 2: 2H2(g) + O2 → 2H2O (l): Expressed in formation of 1 mole: H2(g) + (½)O2 → H2O (l) Equation 3: C2H4(g) + 3O2 → 2CO2(g) + 2H2O (l) Final Equation: 2C (graphite) + 2H2(g) → C2H4(g) Solving for Heat of Formation: Rearrange to create a simultaneous equation so that equations 1 + 2 + 3 = Final equation: 2C +2O2 → 2CO2 (multiplied by 2) = (2)(-393.5) = -787 kJ/mol•C° 2H2 + O2 → 2H2O (multiplied by 2) = (2)(-285.9) = -571.8 kJ/mol•C° 2CO2 + 2H2O → C2H4 + 3O2 (reversed) = 1411.2 kJ/mol•C°

2C + 2H2 → C2H4 52.4 kJ/mol•C° Solving for change in entropy: substitute in values 2C (graphite) + 2H2(g) → C2H4(g) 2(5.7) + 2(130.6) → 219.4 272.6 →219.4 so ΔSf°= -53.2 J/ mol•K Consult the Gibb’s Free Energy Table where ΔHf° > 0 and ΔSf° < 0: Rare: Endergonic reaction, so graphite and hydrogen gas will not combine to form ethylene. Short Cut: The heat of formation of an element in its most stable for (in this example the diatomic oxygen) has a heat of formation of 0 kJ/mol. Establish and balance the combustion equation: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g). The heat of combustion is equal to the ΔHreaction = Hproducts - Hreactants = 2(ΔHwater) + (ΔHcarbon dioxide) - 2(ΔHoxygen) - (ΔHmethane) = (2)(-285.8) + (-392.5) – (2)(0) – (-74.8) = -890.3 kJ.

Heat of Solution Find the heat of solution for KCl dissolved in water. Write the dissolution equation: + - KCl (s) → K (g) + Cl (g): ΔHcrystal lattice = +167.6 kcal Write the solution equation + - + - K (g) + Cl (g) → K (aq) + Cl (aq): ΔHhydration = -163.5 kcal (released heat) Solve for Heat of Solution ΔH = ΔHhydration + ΔHcrystal lattice = -163.5 kcal +167.6 kcal = 4.1 kcal If ΔH > 0, then the reaction is endothermic; if ΔH < 0, then the reaction is exothermic.

La Chatelier’s Principle: Carbon monoxide reacts with hydrogen to form methanol: CO + 2H2 ↔ CH3OH If additional CO is added, the reaction will react to produce additional methanol, dropping the concentration of CO and H2. Nitrogen reacts with hydrogen to form ammonia: N2(g) + 3H2(g) ↔ 2NH3(g) If the volume of the container increases at a constant temperature, then pressure will decrease. A decrease in pressure favors additional moles of gas. Since there are 4 moles gas reactants to 2 moles gas products, the reaction is pushed to the left.

Chemistry Notes

Nuclear Decay and Half-Life Decay Reactions: A A-4 4 4 Alpha Decay: ZX Z-2Y + 2He (the 2He helium nucleus is the alpha particle). A A 0 0 Beta (-) Decay: ZX Z+1Y + -1e (the -1e electron is the beta (-) particle). A A 0 0 Beta (+) Decay: ZX Z-1Y + 1e (the 1e positron is the beta (+) particle). A 0 A Electron Capture: ZX + -1e Z-1Y Gamma Decay: No transmutation occurs Half-Life: Simple Problem: An isotope of cesium (Cs-137) has a half-life of 30 years. If 1.0 mg of Cs-137 disintegrates over a period of 90 years, how many undecayed mg of Cs-137 would remain? Determine the number of half-lifes completed n = t/t0.5 = 90 yr/30 yr = 3 half-lives Determine the remaining mass of Cs-137 n 3 N = N0(½) = (1.0 mg)(½) = 0.125 mg Half-Life: Problem: A 2.5 g sample of an isotope of strontium-90 (Sr-90) was formed in an 1960 explosion of an atomic bomb. The half life of Sr-90 is 28 years. In what year will only 0.625 grams of this isotope remain? n 3 N = N0(½) = (1.0 mg)(½) = 0.125 mg (0.625 g) = (2.5 g)(½)n (0.625 g)/(2.5 g) = (½)n ln(0.25) = ln((½)n) ln(0.25) = n(ln(½)) n = (ln(0.25) / (ln(½)) n = 2 half lives or 56 years

Reduction and Oxidation Reactions: Full: H2 + F2 → 2HF. - - Reduction: F2 + 2e → 2F + - Oxidation: H2 → 2H + 2e Full: Cr2O3(s) + Al(s) → Cr(s) + Al2O3(s) Cr2O3(s): O is -2 so Cr must be +3. Al2O3(s): O is -2 so Al must be +3 Reduction: Cr+3 + 3e- → Cr. +3 - +3 - To balance: 2[Cr + 3e → Cr] so Cr2 + 6e → 2Cr. - + To account for electron and oxygen passage in aqueous solution: Cr2O3 + 6e + 6H → 2Cr + 3H2O +3 - +3 - +3 - Oxidation: Al → Al + 3e . To balance: 2[Al → Al + 3e ] so 2Al → Al2 + 6e - + To account for electron and oxygen passage in aqueous solution: 2Al + 3H2O → Al2O3 +6e + 6H Electrolysis of Pure Water: + - - - Ox: 2H2O(l) → 4H (aq) + O2(g) + 4e . Re: 2H2O(l) + 2e → H2(g) + 2OH (aq). + - - - Balanced: Ox: 2H2O(l) → 4H (aq) + O2(g) + 4e . Re: 4H2O(l) + 4e → 2H2(g) + 4OH (aq). - + - - Combined: 4H2O(l) +2H2O(l) + 4e → 4H (aq) + O2(g) + 4e + 2H2(g) + 4OH (aq). - - + - Simplified: 6H2O(l) + 4e → O2(g) + 4e + 2H2(g) + 4(H + OH )(aq) Eliminate balanced elements and combine: 6H2O(l) → O2(g) + 2H2(g) + 4H2O(l) Cancelled extra waters: 2H2O(l) → O2(g) + 2H2(g)

Solubility Product Constant Solid lead chloride is allowed to dissolve in pure water until equilibrium has been reached and the solution is saturated. The 2+ concentration of Pb is 0.016M. What is the solubility product constant for PbCl2? Write the dissolution equation: 2+ - PbCl2 → Pb (aq) + 2Cl (aq) Determine the moles and reaction exponents for the equation: moles and concentration of dissolved chlorine ions is 2x that of lead. Solve for the solubility product constant p q Ksp = [cation concentration] [anion concentration] 1 2 Ksp = [0.016] [0.032] 17

Chemistry Notes

-5 Ksp = 1.63 x 10

Standard Cell Potential

What is the standard cell potential for a voltaic pile of copper and zinc? 2+ - E°cathode = 0.34 V for Cu + 2e → Cu(s) 2+ - E°anode = -0.76 V for Zn + 2e → Zn(s)

E°cell = E°cathode - E°anode E°cell = 0.34 - -0.76 = 1.10 V

Titration 30.0 mL of a 0.150M nitric acid solution is titrated with calcium hydroxide. The initial burette volume is 0.6 mL and the final burette volume is 22.2 mL. What is the molarity of the calcium hydroxide used in the titration? Determine the chemical formula and the equivalents Nitric acid is HNO3 and yields 1 mole of hydrogen ions per mole of acid. Calcium hydroxide is Ca(OH)2 and yields 2 moles of hydroxide ions per mole of base Determine moles of nitric acid. Molarity = moles/volume so moles = Molarity x Volume = (0.150 moles/Liter)(0.030 Liters) = 0.00450 moles Determine number of moles calcium hydroxide needed to react with moles of hydrogen ions 2 moles of hydroxide per mole of calcium hydroxide, so 0.00225 moles of calcium hydroxide needed. Calculate the molarity of the calcium hydroxide [Ca(OH)2] = moles /Liter = 0.00225 / (22.6 – 0.6) = 0.104M van ‘t Hoff Factor How many grams of benzoic acid must be added to 178 g of water to increase the boiling point temperature by 4°C? Benzoic acid is a non-electrolyte (C7H6O2) with a Kb = 0.52°C/molality Identify the value for i Non-electrolytes do not ionize, so i = 1 Determine the moles of solute 7 x 12 + 6 x 1 + 2 x 16 = 122 g/mole for benzoic acid. Rearrange and substitute into equation

ΔTb = Kb (molality) = Kb(moles solute / kg solvent) = Kb((mass solute / molar mass solute) / kg solvent) ΔTb = Kb((X / molar mass solute) / kg solvent) X = ΔTb (molar mass solute)(kg solvent)/Kb X = (4°C)(122 g/mol)(0.178 kg)/ (0.52°C/molality) = 167 g of benzoic acid.

Chemistry Notes

Tables

Acid/Bases – Strong Acids and Bases Strong Acid Strong Base hydrochloric acid HCl lithium hydroxide LiOH hydrobromic acid HBr sodium hydroxide NaOH hydroiodic acid HI potassium hydroxide KOH nitric acid HNO3 calcium hydroxide Ca(OH)2 sulfuric acid H2SO4 strontium hydroxide Sr(OH)2 perchloric acid HClO4 barium hydroxide Ba(OH)2

Colligative Properties Colligative Property Equation for Property X Proportionality Constant ΔX = Xsolution – Xpure solvent v v Vapor Pressure Lowering ΔP = - P solvent (mole fraction)solute Pure solvent vapor pressure

Boiling Point Elevation ΔTb = Kb (molality) Solvent-dependent constant Kb

Melting Point Lowering ΔTm = - Km (molality) Solvent-dependent constant Km

Osmotic Pressure Posmotic = RT(molarity) Gas constant x Temperature

Gibbs Free Energy

Electrochemical Processes Forced/Electrolytic Cell Spontaneous/Electrochemical Cell Anode Oxidation: Electrons Removed Oxidation: Electrons flow out + - Cathode Reduction: Electrons Forced In Reduction: Electrons flow in - +

Heat Flow In (all changes at 1 atm) Temperature Average Velocity Kinetic Energy Intermolecular Bonds ice T < 0°C ↑ ↑ ↑ ↓ (stretching) ice↔water at 0°C N/A N/A N/A ↓↓ (melting) water at 0°C ↑ ↑ ↑ ↓ water 0°C < T < 100°C ↑ ↑ ↑ ↓ water at 100°C N/A N/A N/A ↓↓ (boiling) vapor at 100°C ↑ ↑ ↑ N/A vapor T > 100°C ↑ ↑ ↑ N/A

Kinetic Molecular Theory

davg b/w density of vavg KEavg # collisions w/ # molecular Pressure molecules closed sys. container collisions V ↑ ↑ ↓ no effect no effect ↓ ↓ ↓ N or n ↑ ↓ ↑ no effect no effect ↑ ↑ ↑ T ↑ no effect no effect ↑ ↑ ↑ ↑ ↑

Chemistry Notes

Solubility Effects Effect of ↑ ↓= decrease, 0 = no change, ↑ = increase, ↑↑ = strong increase Gas solute, Liquid Solvent Solid and Liquid solutes Average KE Gas/Liquid Solubility Solubility for Solubility for collisions at endothermic exothermic interface Pressure 0 ↑↑ ↑ 0 0 Temperature ↑ ↑ ↓ ↑ ↓

Solubility Rules + 1) Salts with ammonium cations (NH4 ) or Group 1 cations are soluble in water - - - - 2) Nitrates (NO3 ), acetates (C2H3O2 ), chlorates (ClO3 ) and perchlorates (ClO4 ) are soluble. - - - 2+ 2+ 2+ 3) Cl , Br , I are soluble except when they are with the following cations: Ag , Hg2 and Pb 2- 2+ 2+ + 2+ 2+ 4) Sulfates (SO4 ) are soluble except with the following cations: Ca , Ba , Ag , Hg2 and Pb - + 2+ 2+ 2+ 5) Hydroxides (OH ) are insoluble except with the following cations: ammonium (NH4 ), Group 1, Ca , Ba , and Sr 2- 2- 3- 2- 6) Sulfites (SO3 ), Sulfides (S ), Phosphates (PO4 ) and Carbonates (CO3 ) are insoluble except with Group 1 cations.

SI Units STANDARD UNITS Quantity Unit Name Symbol Amount Mole Mol Length Meter m Mass Gram g Electric Current Ampere A Temperature Kelvin K Time Second S DERIVED UNITS Acceleration Meters per Second Squared m s-2 Area Square Meter m2 Chemical Reaction Rate Molar per Second M s-1 or mol (L s) -1 Electrical Charge Coulomb C (A s) Electrical Potential, Potential Volts V (J A-1 s-1 or W/A) Difference, Voltage Energy and Heat Joules J (kg m2 s-2) Calories cal (4.18 J) Nutritional Calories Calorie (1 kcal) Force Newton N (kg m s-2) Heat (molar) Joule per Mole J mol-1 Heat Capacity and Entropy Joule per Kelvin J K-1 Mass Atomic Mass Unit u (1.6605 x 10-27 kg) Kilogram* Mass Density Grams per Liter g L-1 Molarity molar mol L-1 Molality molal mol kg-1 Power Watt W (kg m2 s-3) Pressure Pascals Pa (kg m-1 s-2) Atmospheres 1 atm = 101.3 kPa Millimeters of Mercury 760 mm-Hg = 1 atm 20

Chemistry Notes

Specific Heat or Heat Capacity Joules per Kilogram Kelvin J kg-1 K-1 Molar Heat Capacity Joules per Mole Kelvin J mol-1 K-1 Surface Tension Newton per Meter N m-1 or kg s-2 Temperature Degrees Celsius ˚C Velocity (Speed) Meters per Second m s-1 Viscosity Pascal Seconds Pa s Volume Cubic Meter m3 Liter L

Chemistry Notes

Diagrams

Biologically Important Hydrocarbons

Labware Diagrams

Chemistry Notes

Lewis Acid/Base Definitions Example of a Lewis Acid/Base where boron trifluoride (acid) accepts a unshared electron pair from ammonia.

Lewis Dot Structures for Various Molecules

Diagram shows the Lewis Dot structure for Hydronium and Water molecules with the hydrogen bonds shown as dotted lines.

Orbital Quantum Filling

Periodic Chart by Sub-shell

Quantum Orbitals

Chemistry Notes

Sigma (σ) and Pi (π) bonds:

Hybridization Rules: 1s22s22p2 can hybridization in different ways depending on the bonds required.

Carbon orbital structure Carbon hybridized (4 atom bonding) Carbon hybridized (3 atom) Carbon hybridized (2 atom)

Subshell Chart – Electron Energy Levels

Titration Curves

Strong acid titrated with Weak Base titrated with Strong base Strong Acid Polyprotic acid titrated with strong base

VSPER Model Diagrams

Chemistry Notes

Example: Altered bond angles: triagonal pyramidal (107°), bent (104.5°)

Chemistry Notes

Historical Figures

Andre-Marie Ampere (1820): Magnetic Fields Anders Jonas Angstrom: Spectroscopy Svante Arrhenius (1880): Acids (H+ ions) and bases (OH- ions) Amedeo Avogadro: Equal volumes of gas contain equal numbers of molecules, Avogadro’s Number

Neil Bartlett: Nobel Gas Compounds Antoine Henry Becquerel: Radioactivity Daniel Bernoulli (1700): Kinetic Molecular Theory and Fluid Dynamics Henry Bessemer: Bessemer Process for Steel Jons Jakob Berzelius (1800): Modern Chemical Notation Neils Bohr: Discrete Frequencies of Atomic Vibrational Energy Ludwig Boltzmann: Energy Distribution in Gases Robert Boyle (1650): Founder of Modern Chemistry; Boyle’s Law Johannes Bronsted (1920): Bronsted-Lowry definition of acid/base Robert Brown: Brownian Motion

Sadi Carnot (1820): Heat Engines Wallace Carothers: Organic Polymers Jacque Charles (1770): Gas Laws Rudolf Clausius (1860): Entropy Charles Augustin de Coulomb: Coulomb’s Law of Electrical Force Marie Curie: Radioactivity

John Dalton (1800): Indivisible Atom; atoms of element are same; different elements have different atoms Humphry Davy: Electrolysis of Salt Solutions Louis DeBroglie: Wave Characteristics of Mass Democritus: Matter composed of indivisible and indestructible atoms.

Albert Einstein: Energy and Matter Equivalence

Michael Faraday (1830): Induction

Luigi Galvani: Biochemistry Joseph Louis Gay-Lussac: Gas Laws (Law of combining volumes) Hans Wilhelm Geiger: Geiger Counter Josiah Willard Gibbs (1870): Gibbs Free Energy Thomas Graham (1830): Graham’s Law of Effusion and Diffusion (colloid chemistry)

Fritz Haber: Synthesizing Ammonia (Haber Process) Otto Hahn: Nuclear Fission Werner Heisenberg: Uncertainty William Henry: Henry’s Law of gas solubility in liquids Germain Henri Hess: Hess’s Law – Thermodynamics Freidrich Hund: Ground state of atoms (Hund’s Rule)

James Prescott Joule: Mechanical Equivalency of Heat

Irving Langmuir (1920): Surface Chemistry Antoine Lavoisier: Father of modern chemistry: controlled experiments; caloric theory Henri Louis Le Chatelier: La Chatelier’s Principle (changes in chemical equilibrium) Gilbert N. Lewis (1920): Lewis acid/base definition Fritz London: Electrical nature of chemical bonding 26

Chemistry Notes

Thomas Lowry (1920): Bronsted-Lowry definition of acid/base

James Clerk Maxwell: Maxwell’s Equations (propagation of EM waves) Lise Meitner: Nuclear Fission Dmitri Mendeleev: Periodic Table Robert Andrew Millikan: Electron Charge

Sir Isaac Newton: Calculus, Physics, Optics, etc Alfred Nobel: Explosives

Blaise Pascal: (1650): Barometers Louis Pasteur: Chirality Wolfgang Pauli: Pauli’s Exclusion Principle. Linus Pauling (1930): Covalent Bonds Max Planck: Radiation Transfers Energy in Discrete Multiples

William Ramsey: Noble Gases, Alpha Particles Francois Marie Raoult (1870): Colligative properties, Raoult’s Law Wilhelm Konrad Rontgen: x-rays Ernest Rutherford: Gold Foil; mass concentrated in nucleus.

Erwin Schrodinger: Orbitals s, p, d, f Glenn Theodore Seaborg: Synthesis of large atoms; reorganized periodic table Frederick Soddy: Alpha Particles John William Strutt (Lord Rayleigh): Noble Gases

JJ Thomson: charge to mass ratio, plum pudding model William Thomson (Lord Kelvin) (1840): Absolute Temperature

Johannes van der Waals (1870): Intermolecular Attractive Forces Jacobus van’t Hoff (1870): Stereoisomerism of molecules and colligative properties Alessandro Volta: Battery Hermann von Helmsholtz (1840): Conservation of Energy Justus von Liebig: Father of Agricultural Chemistry Friedrich August Kekule von Stradonitz: Benzene Rings

James Watt (1760): Steam Engine Friedrich Wohler: synthesis of an organic compound

Chemistry Notes

Applications

Agriculture and Commerce Fertilizers: chemicals given to promote growth including nitrogen, phosphorus and potassium. A major breakthrough occurred with the development of the Haber Process for Ammonia Production: N2(g) + 3H2(g) ↔ 2NH3 (g) (Fe catalyst) Landfill Processes: Stage 1 – Aerobic Degradation – biodegradable solids react with oxygen to form carbon dioxide and water (releasing heat). Stage 2 – Anaerobic Degradation – Microorganisms that do not require oxygen break down wastes into hydrogen, ammonia, carbon dioxide and inorganic acids. Stage 3 – Production of Methane – When sufficient water and heat are available, microorganisms form gases (carbon dioxide/methane). Pesticides: chemicals given to control or kill organisms. Some pesticides are as harmful to humans as they are to their target. DDT, an insecticide used in the 40’s and 50’s, eradicated malaria but caused widespread birth defects. Roundup, a herbicide used to kill weeds, breaks down quickly in the environment, making it relatively safe.

Medical Antacids: Bases used to control or reduce stomach acids. Antibiotics: organic chemicals that kill or slow the growth of bacteria. Genetic Engineering: Design recombinant DNA to produce human protein molecules. Nuclear Isotopes: X-ray diagnostics and radiology utilize radioactive isotopes. Cobalt-60 is commonly used for external radiotherapy but has been replaced (except in food irradiation). Iodine-131 is used for thyroid treatment.

Nutrition and Cooking Boiling Point Elevation: Adding salt to water increases the boiling point, meaning that boiling water is actually at a higher temperature and therefore cooks food faster. Hydrogenation: Conversion of polyunsaturated oils (with cis- bonds) through hydrogenation into saturated fats to increase the melting point. However, incomplete hydrogenation can create trans-fats, which are even worse than saturated fats for the body. Irradiation: Gamma Rays (Co-60 or Cs-137) are used to kill microorganisms. Nutritional Calories: A unit equivalent to a kilocalorie or 4816 Joules of energy. Oxygen Conversion: Inhaled oxygen is converted to carbon dioxide and water obtaining the same heat of combustion as burning fuel. Preservatives: Nitrites and Nitrates to prevent growth of microorganisms.

Nuclear Power Energy Generation: Conversion of U-235 or Pu-239 through a fission processes causes the release of heat. The heat is used to boil water; the boiling water is harnessed to turn turbines, which generate electricity. Productive fission occurs when a chain reaction is initiated by neutron bombardment of the unstable element.

Chemistry Notes

Molecular Nomenclature

Acids Category Name Example Acids with a hydrogen cation and a hydro ____ic acid hydroiodic acid (HI) monatomic anion hydrotelleric acid (H2Te) Acids with a hydrogen cation and a per______ic acid persulfuric acid (H2SO5) polyatomic anion in its most oxygen rich carbonic acid (H2CO4) state (per____ate)

Acids with hydrogen cation and a ______ic acid sulfuric acid (H2SO4) polyatomic anion in its second most carbonic acid (H2CO3) oxygen rich state (____ate)

Acids with hydrogen cation and a ______ous acid sulfurous acid (H2SO3) polyatomic anion in its second least carbonous acid (H2CO2) oxygen rich state (____ite)

Acids with hydrogen cation and a hypo______ous acid hyposulfurous acid (H2SO2) polyatomic anion in its least oxygen rich state (hypo____ite)

Ions with Unique Names: General Ion Rules: - Thio- means a sulfur is added to the ion; Bi- means hydrogen plus ion (HSO3) is bisulfate; Di- means two of the ion with usually one oxygen lost to bond the ions together.

Monatomic Ions: 2- 2- 2- C2 carbide O2 peroxide S2 disulfide, 1- 1- N3 azide O3 ozonide

Polyatomic Ions: 1- 1- 2- (CH3CO2) acetate (OCN) cyanate (COO)2 oxalate 1- 2- 2- (C8H11N2O3) barbital (Cr2O7) dichromate (C8H4O4) phthalate 1- 1- 2- (C7H5O2) benzoate (HCO2) formate (C4H4O6) tartrate (CN)1- cyanide (OH)1- hydroxide (SCN)1- thiocynate

Oxoanions – Polyatomic Ions with Oxygen: Element hypo____ite _____ite _____ate per____ate 3- 3- As (VA) arsenite (AsO3) arsenate (AsO4) 3- 3- 3- 3- B (IIIA) hypoborite (BO) borite (BO2) borate (BO3) perborate (BO4) - Br (VIIA) bromite (BrO2) 2- 2- 2- C (IVA) carbonite (CO2) carbonate (CO3) percarbonate (CO4) - - - - Cl (VIIA) hypochlorite (ClO) chlorite (ClO2) chlorate (ClO3) perchlorate (ClO4) 2- 2- Cr (VIB) chromate (CrO4) perchromate (CrO5) - - - - I (VIIA) hypoiodite (IO) iodite (IO2) iodate (IO3) periodate (IO4) - - - - N (VA) hyponitrite (NO) nitrite (NO2) nitrate (NO3) pernitrate (NO4) 3- 3- 3- P (VA) hypophosphite (PO2) phosphite (PO3) phosphate (PO4) 2- 2- 2- 2- S (VIA) hyposulfite (SO2) sulfite (SO3) sulfate (SO4) persulfate (SO5) 2- Si (IVA) silicate (SiO3)

Naming rules for ionic and covalents (except organics): 1) Ionic: Cations first, anion last. Covalent: covalent molecules are given in order of increasing electronegativity. 2) For covalent: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca

Chemistry Notes

Naming Rules for organics:

#Carbons Standard Nomenclature Acyl or Nitrile Nomenclature 1 meth form 2 eth acet 3 prop propion 4 but butyr 5 pent pent 6 hex 7 hept 8 oct 9 nono 10 dec

Aromatics: phenyl: C6H5 benzyl: C6H5CH2 (methyl benzene): methyl is toluene, dimethyl is xylene

Hydrocarbon Diagrams Branched Hydrocarbons are named for longest chain, branches ordered alphabetically and using lowest possible numbering system.

Ring bonded hydrocarbons are named cyclo-. Hydrocarbons with double bonds are –ene; triple bonds are –yne.

Hydrocarbons may be named using a variety of rules.

Hydrocarbon Families

Class Functional Structure Suffix Example Group Alcohol Hydroxyl -ol

−OH Primary Secondary ethanol Aldehyde Carbonyl -al

−COH

1 propanal or propionaldehyde

Chemistry Notes

Acid Carboxyl -oic acid

−COOH

methanoic acid or formic acid Ketone Carbonyl -one

2 proponone or acetone Ether Oxy -oxy

\ / or O methyl ethyl ether or ether meth-oxy-ethane Ester oxycarbonyl -yl -oate

(acyl) O | -C-O-

methyl butanoate or methyl butynate Acid carbonyl- -oic (anhydride) oxycarbon anhydride acyl propanoic anhydride or propionoic anhydride Nitrile -C=N -nitrile

ethanonitrile or acetonitrile (acyl) cyanide Amine -N- -amine |

primary secondary tertiary

ethanamine alkyl halide -X fluoro-

chloro- (F, Cl, Br, I) bromo- iodo-

2-bromobutane amide amino -amide carbonyl (acyl)

primary propanamide or propionamide

Chemistry Notes

Hydrocarbon (Organic) Reactions Addition: In many of the organic reactions in which a -bond is replaced with a -bond, the first step is the “attack of a proton by the -bond electron, causing a proton transfer across the double bond. This is followed by the addition of the nucleophile to the remaining cation.

Addition Reactions shown: Hydrochlorination: Ethene + Hydrochloric acid  chloroethane Hydration: Ethene + Water  ethanol Hydrogeneration: Ethene + Hydrogen  Ethane

Substitution: SN2 is a nucleophilic substitution in which a halide replaces hydrogen on a hydrocarbon molecule with a simultaneous attack of the Nucleophile on the central carbon and a elimination of a leaving group. SN1 is a multi-step reaction where the leaving group is eliminated followed by an attack on the cation by a nucleophile with a proton transfer.

Substitution Reactions: - - SN2: OH + CH3Br  CH3OH + Br

Chemistry Notes

Lab Techniques

Hazard Categories Organic Peroxides (R-O-O-R’ ): Very unstable and may self-react to heat, impact, friction, light, or vibration. Organic peroxides should not be used in high school chemistry labs. Peroxide-forming compounds are chemicals that can form organic peroxides through reaction with oxygen in the air if the chemical becomes highly concentrated due to over-boiling or evaporation. Peroxide- forming compounds should be treated as an explosive hazard if they become discolored, layered or crystallize. At-Risk Materials: Ethers (including diipropyl ether, ethyl ether and methyl ether). Potassium metal, Tetrohydrofuran (THF), Cyclohexene, Cyclohexanol, Dioxanes. Lower-Risk: DO NOT BOIL: Isopropanol, 2-butanol, any secondary alcohol Air and Water Reactive Chemicals: The following chemicals react easily with air or water and should not be used in high school chemistry labs. At-Risk Materials: Picric acid, Sodium Metal, Phosphorus Corrosive Materials: The following chemicals are highly corrosive and should not be used in high school chemistry labs. At-Risk Materials: Hydrofluoric Acid (even dilute), Perchloric Acid, Bromine Highly Toxic Materials: The following chemicals are highly toxic and should not be used in high school chemistry labs. At-Risk Materials: Carbon Disulfide, Cyanide compounds, Benzene, Toulene, Mercury and Mercury compounds, Cadmium, Chromium, and Arsenic compounds, carbon tetrachloride and chloroform.

Heating Boiling: Boiling can be made smoother by the addition of boiling stones. Boiling Points can be determined by heating a liquid (with a boiling stone) in a test tube, with a thermometer clamped just above the liquid. The vapor leaving the boiling liquid will be at the same temperature as the boiling liquid. Equipment: Use of Bunsen Burner (high boiling liquids, water and non-flammable liquids) or Hot Plate (containers with flat bottoms) to heat. Melting: Place a pulverized solid in a capillary tube and attach the capillary to a thermometer. This assemble is inserted into a Thiele tube (filled with mineral or silicon oil). Heat the oil using the sidearm on the Thiele Tube; the heated oil delivers an even flow of heat.

Separtation and Filtration Decanting: Separation of a coarse solid from a liquid. Heavy particulates that settle to the bottom of a flask can be removed by carefully pouring out the liquid. Distillation: Separation of two liquids. Liquids with different boiling points can be separated by boiling the mixture and trapping the escaping vapor of the liquid with the lower boiling point. Vapor escapes through the distillation head and travels through a condenser and collects in the receiving flask. Extraction: Separation of two liquids. Two immiscible solutions are poured into a separtory funnel and allowed to settle after being shaken. The lower solution can be removed through the stopcock valve. Gravity Filtration: Separation of solid particulates from a liquid. Remove solids from a liquid by placing folded filter paper in a filtration funnel over a flask. The filtration paper should be wetted with the liquid to seal it to the funnel. Vacuum Filtration: Separation of liquids from a solid. Remove liquids from a solid to obtain a dry solid by using an aspirator or vacuum pump attached to a filter flask (Erlenmeyer flask with a sidearm). This process usually uses a Buchner or Hirsch Funnel.

Storage Rules Chemicals should be stored according to hazard class. Store chemicals away from direct sunlight or localized heat. All chemical containers should be appropriately labeled, and dated upon receipt and opening. Store hazardous chemicals below shoulder height of the shortest person working in the lab. Shelves should be painted or covered with chemical-resistant paint or chemical-resistant coating. Shelves should be secure and strong enough to hold chemicals being stored on them. Do not overload shelves. Personnel should be aware of the hazards associated with all hazardous materials. Separate solids from liquids.

Chemistry Notes

Practice Multiple Choice Questions

1. A piston compresses a gas at a constant temperature. Which gas properties increase? I. Average speed of molecules II. Pressure III Molecular Collisions with container walls per second A. I and II B. I and III C. II and III D I, II, and III

Answer: C Reasoning: In an isothermal reaction, decreasing volume an increase in pressure (Boyle’s Law). Molecular velocity is directly related to temperature, so isothermal means no change in average speed. Since molecules are moving as fast, in a smaller space, they would collide with the container more often. So II and III are true and I is not.

2. The temperature of a liquid is raised at atmospheric pressure. Which liquid property increases? A. Critical Pressure B. Vapor Pressure C. Surface Tension D. Viscosity

Answer: B Reasoning: The critical pressure is the vapor pressure at the critical temperature; above the critical temperature a liquid cannot be formed by an increase in pressure, but with enough pressure a solid may be formed. This is a constant. The vapor pressure of any substance increases non-linearly with temperature. Surface tension and viscosity decrease with temperature.

3. Potassium crystallizes with two atoms contained in each unit cell. What is the mass of potassium found in a lattice 1.00 x 106 unit cells wide, 2.00 x 106 unit cells high and 5.00 x 105 unit cells deep. A. 85.0 ng B. 32.5 g C. 64.9 g D. 130 g

Answer: D Reasoning: Each atom of potassium has a mass of 39.1 u or a molar mass of 39.1 g/mol. The lattice structure contains 2 x 1.00 x 1018 atoms or 0.00000332 moles. At 39.1 g per mole, this is equal to 129.8 x 10-6 g or 130 g

4. A gas is heated in a sealed container. Which of the following occur? A. Gas Pressure rises B. Gas Density decreases C. Average distance b/w molecules increases D. all of the above

Answer: A Reasoning: The system is isovolumetric, so the spacing between molecules and therefore their density is fixed. If heat is added without expansion, then temperature must rise and pressure will rise.

5. How many molecules are in 2.20 pg of a protein with a molecular weight of 150. kDa? A. 8.83 x 109 B. 1.82 x 109 C. 8.83 x 106 D. 1.82 x 106

Answer: C Reasoning: a Da is an acceptable abbreviation of the atomic mass unit which is equivalent to 1.6605 x 10-27 kg or 1.6605 x 10-24 g. A pg is a picogram or 10-12 g. 2.2 x 10-12 g / (150 x 103 Da x 1.6605 x 10-24g/Da) = 8.83 x106 molecules

16 6. At STP, 20. L of O2 contains 5.4 x 10 molecules. According to Avogadro’s hypothesis, how many molecules are in 20. L of Ne? A. 5.4 x 1016 B. 1.0 x 1016 C. 2.7 x 1016 D. 5.4 x 1016

Answer: D The volume of a gas at STP is independent of the atom and depends on the moles of the atom: same number of moles, same volume.

7. An ideal gas at 50.0 ºC and 3.00 atm is in a 300. cm3 cylinder. The cylinder volume changes by moving a piston until the gas is at 50.0 ºC and 1.0 atm. What is the final volume? A. 100. cm3 B. 450. cm3 C. 900. cm3 D. 1.20. dm3

Answer: C

Chemistry Notes

Reasoning: This is an isothermal system since the temperature remains unchanged. Therefore it follows Boyle’s Law. If pressure decreases then volume must increase. (3 atm)(300 cm3) = (1 atm)(V) so V = 900 cm3

8. Which gas law may be used to solve the previous question? A. Charles’s Law B. Boyle’s Law C. Graham’s Law D. Avogadro’s Law

Answer: B Reasoning: Boyle’s law relates pressure and volume in an isothermal system.

9. A blimp is filled with 5000. m³ of helium at 28.0ºC and 99.7 kPa. What is the mass of helium used? A. 797 kg B. 810. kg C. 879 kg D. 8.57 x 103 kg

Answer: A Reasoning: Using the ideal gas law (PV = nRT) and solving for moles: n=PV/RT = (99.7 x 103 Pa)( 5000. m³)/(8.31 Pa m³ mol-1 K-1)(28 +273 K) = 199295.5691 moles of Helium. The molar mass of helium is 4 g/mol so 7.97 x 105 g or 797 kg.

10. Which of the following are able to flow from one place to another? I. Gases II. Liquids III. Solids IV. Supercritical Fluids A. I and II B. II only C. I, II, and IV D. I, II, III, and IV

Answer: C

11. One mole of an ideal gas at STP occupies 22.4 L. At what temperature will one mole of an ideal gas at 1 atm occupy 31.0 L? A. 34.6ºC B. 105 ºC C. 378 ºC D. 442 ºC

Answer: B Reasoning: This problem can be solved using Charles’s Law since the system is isobaric and the number of moles is constant. 22.4 L/273 K = 31.0 L/ X. X = 377.8 K or 105 ºC

12. Why does CaCl2 have a higher normal melting point than NH3? A. London Dispersion Forces in CaCl2 are stronger than covalent bonds in NH3 B. Covalent bonds in NH3 are stronger than dipole-dipole bonds in CaCl2 C. Ionic bonds in CaCl2 are stronger than London Dispersion Forces in NH3 D. Ionic bonds in CaCl2 are stronger than Hydrogen Bonds in NH3

Answer: D Reasoning: CaCl2 is an ionic bond and NH3 is a covalent bond. Weaker intermolecular bonds lead to lower melting points. Covalent bonds are primarily intramolecular bonds. London Dispersion Forces are weaker than covalent bonds. The hydrogen bonds are the intermolecular bonds holding the ammonia together and are weaker than the ionic bonds holding the calcium chloride together.

13. Which intermolecular attraction explains the following trend in straight-chain alkanes? Condensed Structural Formula Boiling Point (ºC) CH4 -161.5 CH3CH3 -88.6 CH3CH3CH3 -42.1 CH3CH3CH3CH3 -0.5 CH3CH3CH3CH3CH3 36.0 CH3CH3CH3CH3CH3CH3 68.7 A. London Dispersion Forces B. Dipole-Dipole Interactions C. Hydrogen Bonding D. Ion-induced Dipole Interactions

Answer: A Reasoning: Alkanes are formed entirely from non-polar C-C and C-H bonds which means there are no dipole interactions or hydrogen bonds. This only leaves the temporary dipoles formed by the electron positions in the atoms.

Chemistry Notes

14. List the substances NH3, PH3, MgCl2, Ne and N2 in order of increasing melting point. A. N2 < Ne < PH3 < NH3 < MgCl2 B. N2 < NH3 < Ne < MgCl2 < PH3 C. Ne < N2 < NH3 < PH3 < MgCl2 D. Ne < N2 < PH3 < NH3 < MgCl2

Answer: D Reasoning: Stronger intermolecular bonds lead to higher melting points. Typically substances with ionic bonds have the highest melting point, MgCl2 is the only ionic substance on the list. Both NH3 and PH3 are polar-covalent and have some dipole-interactions, but NH3 has a greater Electronegativity difference and therefore a slightly greater dipole interaction. Nitrogen has a non-polar covalent bond for a non-spherical structure so would only have weak London dispersion forces. Neon has no intermolecular bonds (very weak London Dispersion Forces and a spherical structure) so it must have the lowest melting point.

15. 1-butanol, ethanol, methanol, and 1-propanol are all liquids at room temperature. Rank them in order of increasing viscosity. A. 1-butanol < 1-propanol < ethanol < methanol B. methanol < ethanol < 1-propanol < 1-butanol C. methanol < ethanol < 1-butanol < 1-propanol D. 1-propanol < 1-butanol < ethanol < methanol

Answer:B Reasoning: Viscosity is internal friction in the movement of the molecules and is due to the intermolecular forces. The larger the molecule the greater the viscosity due to the greater intermolecular forces.

16. Which gas had a diffusion rate of 25% the rate for hydrogen? A. Helium B. Methane C. Nitrogen D. Oxygen

Answer: D. Reasoning: Diffusion is inversely related to the square root of the mass of the molecule. H2 = 2 u He = 4 u, CH4 = 16 u, N2 = 28 u and O2 = 32 u. (Note that Hydrogen, Nitrogen and Oxygen are diatomic). If we set the rate of diffusion for hydrogen as 1 and hold temperature constant, we are looking for a molecule that has the square of 4X the mass of hydrogen: 2 u x 16 = 32 u, or oxygen.

17. 2.00 L of a unknown gas at 1500. mm-Hg and a temperature of 25.0ºC has a mass of 7.52 g. Assuming the ideal gas equation, what is the molecular mass of the gas? A. 21.6 u B. 23.3 u C. 46.6 u D. 93.2 u

Answer: C Reasoning: First converting to standard units 1500 mm-Hg/760 mm-Hg per atm = 1.9737 atm. Solving for number of moles: n = PV/RT = (1.9737 atm)(2.00 L)/(0.08206 atm L mol-1 K-1)(25 + 273 K) = 0.1614 moles. Molar mass is then 7.52 g / 0.1614 = 46.586 u or 46.6 u.

18. Which substance is most likely to be a gas at room temperature? A. SeO2 B. F2 C. CaCl2 D. I2

Answer: B Reasoning: Lowest boiling point substance will be most likely to be a gas at STP. The ionic bonds in calcium chloride are strong, The other molecules are covalent. SeO2 is polar, so will have stronger intermolecular forces than the other two, which are non-polar. The non-polar that is the smallest (smallest London dispersion forces) is F2.

19. What pressure is exerted by a mixture of 2.7 g of hydrogen and 59 g of xenon at STP on a 50. L container? A. 0.69 atm B. 0.76 atm C. 0.80 atm D. 0.97 atm

Answer: C Reasoning: Using Dalton’s Law of partial pressure, the total pressure is the combination of the two pressures. We can find the individual pressures and add. Hydrogen has a molar mass of 2.0 (diatomic) g/mol, so we have 2.7 g / 2 g/mol = 1.35 moles of 36

Chemistry Notes hydrogen. This gives the partial pressure of hydrogen in the 50. L container equal to P = nRT/V = (1.35 mol)(0.08206 atm L mol-1 K- 1)(273)/(50 L) = 0.605 atm. The moles of xenon with a molar mass of 131 g/mole is n = 59 g/131 g/mol = 0.450 moles, and a partial pressure of P = nRT/V = (10.450 mol)(0.08206 atm L mol-1 K-1)(273)/(50 L) = 0.202 atm. The total pressure would be 0.605 atm + 0.202 atm = 0.807.

20. A few minutes after opening a bottle of perfume, the scent is detected on the other side of the room. What law relates to this phenomenon? A. Graham’s Law B. Dalton’s Law C. Boyle’s Law D. Avogadro’s Law

Answer: A Reasoning: Graham’s law covers the diffusion of gas molecules.

21. Which of the following are true? A. Solids have no vapor pressure B. Dissolving a solute in a liquid increases its vapor pressure C. The vapor pressure of a pure substance is characteristic of that substance and of its temperature D. All of the above

Answer: C Reasoning: Solids do have a vapor pressure. Dissolving a solute in a liquid causes the vapor pressure to decrease.

22. Find the partial pressure of N2 in a container at 150 kPa holding H2O and N2 at 50ºC. The vapor pressure of H2O at 50ºC is 12 kPa. A. 12 kPa B. 138 kPa C. 162 kPa D. The value cannot be determined

Answer: B Reasoning: Dalton’s law of partial pressure says that the pressure of the system is equal to the sum of the pressure of the individual components in the system. If the vapor pressure of H2O is 12 kPa, then the pressure of the nitrogen is 150 kPa – 12 kPa = 138 kPa

23, The normal boiling point of water on the Kelvin scale is closest to: A. 112 K B. 212 K C. 273 K D. 373 K

Answer: D Reasoning: The normal boiling point of water is given as 100ºC, converting Celsius to Kelvin, Tbp = 373 K

24. Which phase may be present at the triple point of a substance? I. Gas, II. Liquid, III. Solid, IV. Supercritical Fluid A. I, II, and III B. I, II and IV C. II, III, and IV D. I, II, III, and IV

Answer: A Reasoning: The supercritical fluid occurs at the critical point, not the triple point. The triple point is where the phases (solid, liquid and gas) can all exist.

25. In the following phase diagram, ______occurs as P is decreased from A to B at a constant T and ______occurs as T is increased from C to D at constant P. A. deposition, melting B. sublimation, melting C. deposition, vaporization D. sublimation, vaporization

Answer: D Reasoning: A to B shows a process going from solid to gas at a constant temperature (sublimation) and C to D shows a process going from liquid to gas at a constant pressure (vaporization).

26. Heat is added to a pure solid at its melting point until it all becomes liquid at its freezing point. Which of the following occur? A. Intermolecular attractions are weakened B. The kinetic energy of the molecules does not change C. The freedom of the molecules to move about increases D. All of the above.

Chemistry Notes

Answer: D Reasoning: During the phase change process, all added energy is going towards changing the intermolecular bonds and does not increase the kinetic energy of the molecules.

27 Which of the following occur when NaCl dissolves in water? A. Heat is required to break bonds in the NaCl crystal lattice B. Heat is released when hydrogen bonds in water are broken C. Heat is required to form the bonds of hydration D. The oxygen end of the water molecule is attracted to the Cl- ion.

Answer: A Reasoning: The oxygen end of a water molecule has a slight negative polarity so the chlorine ion would repel from it. Heat is released during formation of hydration; heat is required to break bonds for both water and sodium chloride.

28. The solubility of CoCl2 is 54 g per 100 g of ethanol. Three flasks each contain 100 g of ethanol. Flask #1 also contains 40 g CoCl2 in solution. Flask #2 contains 56 g CoCl2 in solution. Flask #3 contains 5 g CoCl2 in solid form in equilibrium with 54 g of CoCl2 in solution. Which of the following describe the solutions present in the liquid phase of the flasks? A. #1-saturated, #2-supersaturated, #3-unsaturated B. #1-unsaturated, #2-miscible, #3-saturated C. #1-unsaturated, #2-supersaturated, #3-saturated D. #1-unsaturated, #2-not at equilibrium, #3-miscible

Answer: C Reasoning: Solution #1 has less solute than the solution is capable of holding, therefore it is unsaturated. Solution #2 has more solute than the solution would normally be able to hold, therefore it is supersaturated. Solution #3 is at equilibrium at the maximum saturation so it is saturated.

29. The solubility at 1.0 atm of pure CO2 is water at 25ºC is 0.034M. According to Henry’s Law, what is the solubility at 4.0 atm of pure CO2 in water at 25 ºC? Assume no chemical reaction occurs between CO2 and H2O. A. 0.0085M B. 0.034M C. 0.14M D. 0.25M

Answer: C Reasoning: Henry’s Law states that the solubility of a gas in a liquid at a particular temperature is proportional to the pressure of that gas above the liquid. Therefore, if pressure increases by 4X, then solubility should increase by 4X leading to 4X the concentration.

30. Carbonated water is bottled at 25 ºC under pure CO2 at 4.0 atm. Later the bottle is opened at 4 ºC under air at 1.0 atm that has -4 a partial pressure of 3 x10 atm CO2. Why do CO2 bubbles form when the bottle is opened? A. CO2 falls out of solution due to a drop in solubility at the lower total pressure. B. CO2 falls out of solution due to a drop in solubility at the lower CO2 pressure. C. CO2 falls out of solution due to a drop in solubility at the lower temperature. D. CO2 is formed by the decomposition of carbonic acid.

Answer:B Reasoning: Lowering the pressure above a liquid/gas solution does decrease the solubility but the gas that is important is the gas in the solution. Lowering the temperature increases the solubility and the temperature range in this problem is minimal. Carbonic acid may decompose into carbon dioxide but this would only be a very small fraction of the carbon dioxide released.

31. When KNO3 dissolves in water, the water grows slightly colder. An increase in the temperature will ______the solubility of KNO3. A. Increase B. Decrease C. Have no effect on D. Have an unknown effect with the information given.

Answer: A Reasoning: KNO3 has a positive trend in the relationship between solubility and temperature.

Chemistry Notes

32. An experiment requires 100. mL of a 0.500M solution of MgBr2. How many grams of MgBr2 will be present in solution? A. 9.21 g B. 11.7 g C. 12.4 g D. 15.6 g

Answer: A Reasoning: MgBr2 has a molar mass of 24.31 + 2x79.90 = 184.11 g/mol (184 u). The number of moles is equal to the molarity multiplied by the volume so moles = 0.500 (mol/L)(0.100 L) = 0.05 moles. So the mass is 0.05 moles x 184.11 g/mol = 9.21 g

33. 500 mg of RbOH are added to 500. g of ethanol resulting in 395 mL of solution. Determine the molarity and molality of RbOH. A. 0.0124M, 0.00488m B. 0.0124M, 0.00976m C. 0.0223M, 0.00488m D. 0.0223M, 0.00976m

Answer: B Reasoning: There are 500 mg of RbOH with a molar mass of 85.5 + 16.0 + 1.0 = 102.5 g/mol so there are (0.5 g)/(102.5 g/mol) = 0.00488 moles in the solution. The molarity will be 0.00488 moles / 0.395 L = 0.0124M. The molality will be 0.00488 mol / 0.5 kg = 0.00976m.

34. 20.0 g H3PO4 in 1.5 L of solution are intended to react with KOH according to the following reaction: H3PO4 + 3KOH K3PO4 + 3H2O. What is the molarity and normality of the H3PO4 solution? A. 0.41M, 1.22N B. 0.41M, 0.20 N C. 0.14M, 0.045N D. 0.14M, 0.41N

Answer: D Reasoning: H3PO4 has a molar mass of 3x1.0 + 31.0 + 4x16.0 = 98 g/mole, so 20 g is 0.204 moles H3PO4. The molarity of the solution is 0.204 moles / 1.5 L = 0.14M. Every 1 equivalent of H3PO4 reacts with 3 equivalents of KOH, so the normality would be 3x(0.14) = 0.41N

35. Aluminum sulfate is a strong electrolyte. What is the concentration of all species in a 0.2M solution of aluminum sulfate? 3+ 2- 3+ 2- 3+ 2- A. 0.2M Al , 0.2M SO4 B. 0.4M Al , 0.6M SO4 C. 0.6M Al , 0.2M SO4 D. 0.2M Al2(SO4)3

Answer: B Reasoning: Aluminum has an oxidation number of 3+; sulfate has an oxidation number of 2-, so the molecular formula of Aluminum Sulfate is Al2(SO4)3. When it dissolves into solution as an electrolyte, each mole of aluminum sulfate produces 2 moles of aluminum ions and 3 moles of sulfate ions. This means the molarity of the aluminum is 0.4M and the molarity of the sulfate is 0.6M.

36. 15 g of formaldehyde (CH2O) are dissolved in 100. g of water. Calculate the weight percentage and the mole fraction of formaldehyde in the solution. A. 13%, 0.090 B. 15%, 0.090 C. 13%, 0.083 D. 15%, 0.083

Answer: C Reasoning: The molar mass of formaldehyde is 12+2+16 = 30 g/mol, so 15 g / 30 g/mol = 0.5 moles. Water has a molar mass of 18 g/mole so the solution contains 100 g / 18 g/mol = 5.56 moles water. The weight percentage is 15 g /(100 g + 15 g) = 13%. The mole fraction is 0.5 moles / (0.5 moles + 5.56 moles) = 0.083.

37. Which of the following would make the best solvent for Br2? A. H2O B. CS2 C. NH3 D. molten NaCl

Answer: B Reasoning: “Like dissolves like”; since Br2 is non-polar covalently bonded with London Dispersion Forces, the solvent should be non- polar covalent. NaCl is ionic, H2O and NH3 are both polar covalent with hydrogen bonds, so CS2 would be the best solvent.

38. Which of the following is most likely to dissolve in water? A. H2 B. CCl4 C. (SiO6)n D. CH3OH

Answer: D

Chemistry Notes

Reasoning: H2O has relatively strong hydrogen bonds, so the best solute would also have relatively strong hydrogen bonds. (SiO6)n is a network solid acting as one large molecule. CCl4 is covalent with weak London dispersion forces. H2, since it is polar and there are no relative negatives, is bonded with London Dispersion Forces. CH3OH has hydrogen bonds and will mix with water.

39. Which of the following is not a colligative property? A. Viscosity lowering B. Freezing point lowering C. Boiling point elevation D. Vapor pressure lowering

Answer: A Reasoning: The addition of solutes to a liquid solvent drives equilibrium toward the liquid phase and increasing boiling point and decreasing melting point. This also causes the vapor pressure to lower.

40. BaCl2 (aq) + Na2SO4 (aq) BaSO4 (s) + 2NaCl (aq) is an example of a ______reaction. A. acid-base B. precipitation C. redox D. nuclear

Answer: B Reasoning: The reaction is obviously not nuclear since there is no emission of radiation particles or change in energy state of the nucleus. The oxidation number on the Ba, Cl, Na and SO4 ions remain unchanged so this is not a redox reaction. Neither reactant is an acid or a base, so it is not an acid-base reaction. The aqueous solution falls out into solid form, so this is a precipitation.

41. List the following aqueous solutions in order of increasing boiling point. I. 0.050m AlCl3, II. 0.080m Ba(NO3)2, III. 0.090m NaCl and IV. 0.12m ethylene glycol (C2H6O2). A. I < II < III < IV B. I < III < IV < II C. IV < III < I < II D. IV < III < II < I

Answer: C Reasoning: Increasing boiling point is a colligative property, so the solution with the most particles will have the highest boiling point. Ethylene glycol does not ionize in water so 1 mole forms 1 mole particles or 0.12m. Sodium Chloride ionizes into 2 particles, so 1 mole forms 2 moles particles or 0.18m. Barium Nitrate ionizes into 3 particles (0.24m); Aluminum Chloride ionizes into 4 particles (0.20m). So 0.24m > 0.20m > 0.18m > 0.12 m

42. Osmotic pressure is the pressure required to prevent ______flowing from low to high ______concentration across a semipermeable membrane. A. solute, solute B. solute, solvent C. solvent, solute D. solvent, solvent

Answer: C Reasoning: Solvents flow….concentration is based on solute present.

43. A solution of NaCl is water is heated on a mountain in an open container until it boils at 100ºC. The air pressure on the mountain is 0.92 atm. According to Raoult’s Law, what mole fraction of Na+ and Cl- are present in the solution. A. 0.04 Na+, 0.04 Cl- A. 0.08 Na+ , 0.08 Cl- A. 0.46 Na+ , 0.46 Cl- A. 0.92 Na+, 0.92 Cl-

Answer: A Reasoning: The vapor pressure is the pressure at boiling point, so the vapor pressure of the solution was 0.92 atm. At 0.92 atm, pure water would boil at a temperature less than 100ºC, the lowering in boiling point due to pressure was offset by the boiling point vapor vapor elevation from the addition of solute. Raoult’s Law states P solution = P pure solvent (mole fraction)solvent , so to solve for the mole vapor vapor fraction = P solution / P pure solvent = 0.92 atm at 100 ºC / 1.0 atm at 100 ºC = 0.92 mol H2O/total moles. This leaves a mole fraction of 0.08. Since 1 mole of NaCl dissolves into 1 mole of sodium ions and 1 mole of chlorine ions, the mole fraction for each is 0.04.

44. Write a balanced nuclear equation for the emission of an alpha particle by polonium-209. 209 205 4 209 205 4 209 209 0 209 205 4 A. 84Po 81Pb + 2He B. 84Po 82Bi + 2He C. 84Po 85At + -1e D. 84Po 82Pb + 2He

Answer: D

Chemistry Notes

4 Reasoning: The 2He particle is the alpha particle. The process of alpha decay decreases the mass number by 4 and the atomic number by 2. Therefore Polonium-209, which has an atomic number of 84, will decay to Lead-205, which has an atomic number of 82.

45. Write a balanced nuclear equation for the decay of calcium-45 to scandium-45. 45 41 4 45 0 45 45 45 0 45 0 45 A. 20Ca 18Sc + 2He B. 20Ca + -1e 21Sc C. 20Ca 21Sc + -1e D. 20Ca + 1p 21Sc

Answer: C Reasoning: The decay process that takes Calcium to Scandium must increase the atomic number by 1. In this case, the mass number is also unchanged. This happens through beta decay.

3 3 3 46. 1 H decays with a half-life of 12 years. 3.0 g of pure 1H were placed in a sealed container 24 years ago. How many grams of 1 H remain? A. 0.38 g B. 0.75 g C. 1.5 g D. 3.0 g

Answer: B Reasoning: The hydrogen isotope was in the sealed container for 2 half-lives, therefore it is expected that there are approximately ¼ the original hydrogen isotopes remaining undecayed.

47. Oxygen-15 has a half-life of 122 seconds. What percentage of a sample of oxygen-15 has decayed after 300 seconds? A. 18.2% B. 21.3% C. 78.7% D. 81.8%

Answer: D Reasoning: The oxygen has gone through n = 300/122 = 2.459 half-lives, so the fraction left is 1/22.459 = 0.1819 or 18.2% is remaining and 100-18.2 = 81.8% has already decayed.

48. Which of the following isotopes is commonly used for medical imaging in the diagnosis of diseases? A. cobalt-60 B. technetium-99 C. tin-117 D. plutonium-238

Answer: B

49. Carbon-14 dating would be useful in obtaining the age of which object? A. a 20th century Picasso painting B. a mummy from ancient Egypt C. A dinosaur fossil D. all of the above

Answer: B Reasoning: Carbon-14 has a half-live of 5730 years. Too little carbon will have decayed in the painting and too much in the dinosaur.

50. Which of the following isotopes can create a chain reaction of nuclear fission? A. Uranium-235 B. Uranium-238 C. Plutonium-238 D. All of the above

Answer: A

51. List the following in chronological order from earliest to most recent with respect to their most significant contribution to atomic theory. I. John Dalton, II. Niels Bohr, III. JJ Thomson, IV Ernest Rutherford. A. I, III, II, IV B. I, III, IV, II C, I, IV, III, II D, III, I, II, IV

Answer: B Reasoning: John Dalton’s work was on the individual nature of atoms of each element and he thought the atom as indestructible. JJ Thomson believed that the negative charges in an atom were scattered in its positive mass (plum pudding). Ernest Rutherford determined the concentration of mass and positive charge in an atom to be centralized in the nucleus. Neils Bohr worked in early quantum atomic theory.

Chemistry Notes

52. Match the theory with the scientist who first proposed it: I. Electrons, atoms, and all objects with momentum also exist as waves. II. Electron density may be accurately described by a single mathematical equation. III. There is an inherent indeterminacy in the position and momentum of particles. IV. Radiant energy is transferred between particles in exact multiples of a discrete unit. A. I-deBroglie, II-Planck, III-Schrodinger, IV-Thomson B, I-Dalton, II-Bohr, III-Planck, IV-deBroglie C. I-Henry, II-Bohr, III-Heisenberg, IV- Schrodinger D. I-deBroglie, II- Schrodinger, III-Heisenberg, IV-Planck

Answer: D

53. How many neutrons are in Cobalt-60? A. 27 B. 33 C. 60 D. 87

Answer: B Reasoning: Cobalt-60 has an atomic number of 27 (27 protons) and a mass number of 60 (60 nucleons). The number of neutrons is the number of nucleons less the number of protons, or 33.

54. The terrestrial composition of an element is 50.7% as an isotope with an atomic mass of 78.9 u and 49.3% as an isotope with an atomic mass of 80.9 u. Both isotopes are stable. Calculate the atomic mass of the element. A. 79.0 u B. 79.8 u C. 79.9 u D. 80.8 u

Answer: C Reasoning: The weighted average of the two isotopes is the atomic mass of the element: (0.507)(78.9) + (0.493)(80.9) = 79.9 u

55. Which of the following is a correct electron arrangement for oxygen?

A. B. 1s21p22s22p2 C. 2, 2, 4 D. None of the above

Answer: D Reasoning: Oxygen has 8 electrons. The first diagram violates Hund’s Rule. The second choice shows the 1p orbital which doesn’t exist. The third option shows the simplified orbital level of 2 in the 1s shell, but should then show 6 in the 2sp shell.

56. Which of the following statements about radiant energy is not true? A. The energy change of an electron transition is directly proportional to the wavelength of the emitted or absorbed photon. B. The energy of an electron in a hydrogen atom depends only on the principle quantum number C. The frequency of photons striking a metal determines whether the photoelectric effect will occur. D. The frequency of a wave of electromagnetic radiation is inversely proportional to its wavelength.

Answer: A Reasoning: The energy change of an electron is inversely proportional to the wavelength of the photon (ΔE = hc/λ). The energy of an electron in hydrogen atom depends only on its principle quantum number (but that has nothing to do with radiant energy). The frequency of the photons does determine whether the photoelectric effect will occur (the frequency must be above the threshold frequency). Frequency of any wave is inversely proportional to its wavelength (f = v/ λ).

57. Match the orbital diagram for the ground state of carbon with the rule or principle it violates:

A. I – Pauli Exclusion Principle, II – Aufbau, III – no violation, IV – Hund’s Rule B. I – Aufbau, II – Pauli Exclusion Principle, III – no violation, IV – Hund’s Rule C. I – Hund’s Rule, II – no violation, III – Pauli Exclusion Principle, IV – Aufbau D. I – Hund’s Rule, II – no violation, III – Aufbau, IV – Pauli Exclusion Principle

Answer: C 42

Chemistry Notes

Reasoning: The first diagram fails to fully fill each orbital before adding a second electron to the orbital (Hund’s Rule). The second diagram is correct. The third diagram has two electrons with the identical spin in the 2s orbital (Pauli’s Exclusion Principle) and the fourth diagram fails to fully fill the 2s orbital before going to the 2p orbital (Aufbau).

58. Select the list of atoms that are arranged in order of increasing size. A. Mg, Na, Si, Cl B. Si, Cl, Mg, Na C. Cl, Si, Mg, Na D. Na, Mg, Si, Cl

Answer: C Reasoning: All the atoms shown are in the third period. The trend in atomic size as you go from Group I to Group 18 in a period is to decrease in size (due to the additional electromagnetic interaction between the nucleus and the electrons. Chlorine is Group 17, Silicon is Group 14, Magnesium is Group 2 and Sodium is Group 1.

59. Based on trends in the periodic table, which of the following properties would you expect to be greater for Rb than for K? I. Density, II. Melting Point, III. Ionization Energy, and IV. Oxidation number in a compound with Chlorine. A. I only B. I, II, and III C. II and III D. I, II, III and IV

Answer: A Reasoning: Rb and K are both in Group I; Rb is in period 5 and K is in period 4. Rb will have a greater diameter and greater mass (more particles), a lesser ionization energy (electrons are more removed from the nucleus). They will both have the same oxidation number with Chlorine (group I atoms have an oxidation number of +1 with Group 17 atoms). Melting point for metals decreases with size. (Note: if you realize that III and IV are not true, the only remaining choice is A).

60. Which oxide forms the strongest acid in water? A. Al2O3 B. Cl2O7 C. As2O5 D. CO2

Answer: B Reasoning: Acid strength increases with the electronegativity and with oxidation state. The trend in electronegativity is increasing from lower left to upper right, so Chlorine has the greatest electronegativity. The oxidation states are Al (+3), Cl (+7), As(+5), C(+4). This means that chlorine has the greatest electronegativity and the greatest oxidation state. (The greatest acid releases the greatest amount of hydrogen ions. This occurs when there is the most amount of oxygen available for the formation of hydroxide ions. The dichlorine heptoxide atom has the greatest amount of oxygen.)

61. Rank the following bonds from least to most polar: C-H, C-Cl, H-H, C-F A. C-H < H-H < C-F < C-Cl B. H-H < C-H < C-F < C-Cl C. C-F < C-Cl < C-H < H-H D. H-H < C-H < C-Cl < C-F

Answer: D Reasoning: H-H bonds are non-polar. C-H bonds are also non-polar (but slightly less so). C-Cl is polar but its difference in electronegativity is less than that of C-F.

62. At room temperature, CaBr2 is expected to be: A. ductile solid B. brittle solid C. soft solid D. gas

Answer: B Reasoning: CaBr2 is a ionic salt – brittle solids.

63. Which of the following is a proper Lewis dot structure of CHClO?

Answer: C Reasoning: With the exception of hydrogen, the atoms in the molecule follow the octet rule.

64. In C2H2, each carbon atom contains the following valence orbitals.

A. p only B. p and sp hybrids C. p and sp2 hybrids and D. sp3 hybrids

Chemistry Notes

Answer: B Reasoning: In carbon bonds, the 2s electrons are promoted to hybrid bonds. If carbon bonds to 4 atoms then all the s and p electrons become sp3; in this molecule (ethyne), each carbon has a triple bond to the other carbon and a single bond to a hydrogen. This means that the bond structure is sp and p bonds.

65. Which statement about molecular structures is false?

A. is a conjugated molecule B. A bonding sigma (σ) orbital connects two atoms by the straight line between them. C. A bonding pi (π) orbital connects two atoms in a separate region from the straight line between them.

D. The anion with resonance forms will always exist in one form or the other.

Answer: D Reasoning: The 1, 3 – butene molecule is conjugate since the adjacent middle carbons both contain double bonds. The sigma bond lies on a straight-line axis between two atoms; the pi bond lies off this line. The resonance forms describe the structure but the atom actually exists between the two structures.

66. What is the shape of the PH3 molecule? Use the VSEPR model. A. Trigonal pyramidal B. Trigonal bipyramidal C. Trigonal planar D. Tetrahedral

Answer: A Reasoning: The central potassium atom has an electron configuration of: 1s2 2s2 2p6 3s2 3p3; this means there are 5 valence electrons. The electrons bond to the hydrogen atoms with three bonds and 1 unshared pair. This pushes the hydrogens out of the plane, forming a trigonal pyramidal shape.

67. What is the chemical composition of magnesium nitrate? A. 11.1% Mg, 22.2% N, 66.7% O B. 16.4% Mg, 18.9% N, 64.7% O C. 20.9% Mg, 24.1% N, 55.0% O D. 28.2% Mg, 16.2% N, 55.7% O

Answer: B Reasoning: Magnesium Nitrate is Mg(NO3)2 since Magnesium has an oxidation number of +2 and Nitrate has an oxidation number of -1. Therefore each magnesium nitrate atom has 1 Magnesium atom (24.3 u), 2 Nitrogen atoms (14.0 u each) and 6 Oxygen atoms (16.0 u each). The total atomic mass of Magnesium Nitrate is 24.3 + 2(14.0) + 6(16.0) = 148.3 u. The percentage then is 16.4% Mg, 18.9% N, and 64.7% O.

68. The IUPAC name for Cu2SO3 is A. Dicopper sulfur trioxide B. Copper (II) sulfate C. Copper (I) sulfite D. Copper (II) sulfide.

Answer: C Reasoning: The first choice is improper for ionic/metal to non-metal bonds. Additionally, SO3 is a named polyatomic ion, sulfite, with an oxidation number of -2. Since 1 sulfite is bonding to 2 coppers, the oxidation number of the copper must be +1. Therefore the appropriate name is copper (I) sulfite.

69. Which name or formula is not represented properly? A. Cl4S B. KClO3 C. Calcium dihydrogen phosphate D. Sulfurous Acid.

Answer: A

Chemistry Notes

Reasoning: The potassium chlorate molecule (B) is appropriately written. Sulfurous Acid is also appropriate for the molecule H2SO3. Tetrachlorine sulfide should be given as sulfide tetrachloride since the covalent molecules are usually given in order of increasing electronegativity.

70. Household “chlorine bleach” is sodium hypochlorite. Which of the following best represent the production of sodium hypochlorite, sodium chloride and water by bubbling chlorine gas through aqueous sodium hydroxide? A. 4Cl(g) + 4NaOH(aq)→NaClO2(aq) + 3NaCl(aq) + 2H2O(l) B. 2Cl2(g) + 4NaOH(aq)→NaClO2(aq) + 3NaCl(aq) + 2H2O(l) C. 2Cl(g) + 2NaOH(aq)→NaClO(aq) + NaCl(aq) + H2O(l) D. Cl2(g) + 2NaOH(aq)→NaClO(aq) + NaCl(aq) + H2O(l)

Answer: D 1- Reasoning: Chlorine is diatomic and should be written as Cl2, hypochlorite is ClO , and sodium chloride is NaCl. Even without looking at balancing, the only option is D.

71. Balance the equation for the neutralization reaction between phosphoric acid and calcium hydroxide by filling in the blank stoichiometric coefficients. ___H3PO4 + ___Ca(OH)2 → ___Ca3(PO4)2 + ___H2O A. 4, 3, 1, 4 B. 2, 3, 1, 8 C. 2, 3, 1, 6 D. 2, 1, 1, 2

Answer: C Reasoning: It may be easiest to try each solution rather than attempting to balance. You may also write the half reactions and balance the hydrogen and electron movement or attempt to balance normally. Initially we have reactants products (2x)4 +(3x)2 oxygen = 14 8 + (6x)1 oxygen = 14 (2x)3 + (3x)2 hydrogen = 12 (6x) 2 hydrogen = 12 (2x)1 phosphorous 2 phosphorous (3x)1 calcium 3 calcium 2H3PO4 + 3Ca(OH)2 → 1Ca3(PO4)2 + 6H2O

72. Write an equation showing the reaction between calcium nitrate and lithium sulfate in aqueous solution. Include all products. A. CaNO3(aq) + Li2SO4(aq) →CaSO4(s) + Li2NO3(aq) B. Ca(NO3)2(aq) + Li2SO4(aq) →CaSO4(s) + 2LiNO3(aq) C. Ca(NO3) 2(aq) + Li2SO4(aq) → 2LiNO3(s) + CaSO4(aq) D. Ca(NO3)2(aq) + Li2SO4(aq) + 2H2O(l)→ LiNO3(aq) + CaSO4(s) + H2SO4(aq)

Answer: B Reasoning: Calcium Nitrate is Ca(NO3)2, Lithium Sulfate is Li2SO4; this eliminates A. To determine the possible of precipitation, we must look at solubility. Nitrates are soluble; sulfates are soluble except with Ca, Ba, Ag, Hg, and Pb cations; in a double replacement, where calcium nitrate and lithium sulfate combine, they will form lithium nitrate (aq) and calcium sulfate (s). So there must be a precipitation. Option D also shows a dissolution of the water molecule which would not occur with the addition of these two salts.

73. Find the mass of carbon dioxide produced by the combustion of 15 kg of isopropyl alcohol (C3H7OH) in the reaction: 2C3H7OH + 9O2 → 6CO2 + 8H2O A. 33 kg B. 44 kg C. 50 kg D. 60 kg

Answer: A Reasoning: 2 moles of isopropyl alcohol react with 9 moles of oxygen to form 6 moles of carbon dioxide and 8 moles of water. So 15 kg of isopropyl alcohol (molecular mass of 60 g/mol) is 250 moles. At a 2:6 ratio, then 750 moles of carbon dioxide is formed. Carbon dioxide has a molar mass of 44 g/mol, so forms 33 kg.

74. What is the density of nitrogen gas at STP? Assume ideal gas and a value of 0.08206 L•atm•mol-1•K-1 A. 0.62 g/L B. 1.14 g/L C. 1.25 g/L D. 2.03 g/L

Answer: C Reasoning: 1 mole of any ideal gas at STP has a volume of 22.4 L. 1 mole of nitrogen gas has a mass of 28 g. Therefore, the density is 28 g / 22.4 L = 1.25 g/L

Chemistry Notes

75. Find the volume of methane that will produce 12 cubic meters of hydrogen in the reaction: CH4(g) + H2O(g) → CO(g) + 3H2(g). Assume temperature and pressure are constant. A. 4.0 m³ B. 32 m³ C. 36 m³ D. 64 m³

Answer: A Reasoning: 1 mole of methane reacts with 1 mole of water to form 1 mole of carbon monoxide and 3 moles of hydrogen. So proportionally, volume must be 4 cubic meters.

76. A 100. L vessel of pure oxygen gas at 500. kPa and 20°C is used for the combustion of butane: 2C4H10 + 13O2 → 8CO2 + 10H2O. Find the mass of butane to consume all the oxygen gas in the vessel. Assume oxygen gas is an ideal gas. A. 183 g B. 467 g C. 1.83 kg D. 7.75 kg

Answer: A Reasoning: Finding the moles of oxygen: n = PV/RT = (500,000 Pa)(0.100 m³) / (8.31 Pa•m³•mol-1•K-1)(20 + 273 K) = 20.54 moles. Since oxygen to butane is 13:2 ratio, then 20.54 oxygen requires 3.16 moles. The molar mass of butane is 58 g/mole, so the mass of butane must be 183 g.

77. Consider the reaction between iron and hydrogen chloride gas: Fe(s) + 2HCl (g) → FeCl2(s) + H2(g). 7 moles of iron and 10 moles of HCl react until the limiting reagent is consumed. Which statements are true? I. HCl is the excess reagent. II. HCl is the limiting reagent. III. 7 moles of H2 are produced. IV. 2 moles of the excess reagent will remain. A. I and III B. I and IV C. II and III D. II and IV

Answer: D Reasoning: 1 mole of iron reacts with 2 moles of hydrochloric acid (hydrogen chloride gas). 10 moles of hydrochloride gas would react completely with 5 moles of iron. So HCl is the limiting reagent.

78. 32.0 g of hydrogen and 32.0 g of oxygen react to form water until the limiting reagent is consumed. What is present in the vessel after the reaction is complete? A. 16.0 g O2, 48.0 g H2O B. 24.0 g H2, 40.0 g H2O C. 28.0 g H2, 36.0 g H2O D. 28.0 g O2, 34.0 g H2O

Answer: C Reasoning: The equation for this reaction is 2H2 + O2 → 2H2O , so two moles of hydrogen are needed for every one mole of oxygen. 32 g / 2 g/mol = 16 moles hydrogen and 32 g / 32 g/mol = 1 mole oxygen. So, since the oxygen is the limiting reagent, we will consume only 2 moles of hydrogen and produce 2 moles of water. The remaining 14 moles of hydrogen have a mass of 28 g, and the 2 moles of water have a mass of 36 g.

79. Three experiments were performed at the same initial temperature and pressure to determine the rate of the reaction: 2ClO2(g) + F2(g) → 2ClO2F(g). Results are shown in the table below. Concentrations are given in millimoles per liter. What is the rate law of this reaction?

Experiment Initial [ClO2] in mM Initial [F2] in mM Initial rate of [ClO2F] increase 1 5.0 5.0 0.63 2 5.0 20 2.5 3 10 10 2.5 2 2 A. Rate = k[F2] B. Rate = k[ClO2][F2] C. Rate = k[ClO2] [F2] D. Rate = k[ClO2][F2]

Answer: B Reasoning: The rate law equation is given as rate = k[reactant 1]a[reactant 2]b. When fluorine concentration was quadrupled, the initial production rate quadrupled, when chlorine dioxide is held constant. When chlorine dioxide is doubled and fluorine is reduced to half, the initial rate is unchanged.

- - 80. The reaction (CH3)CBr(aq) + OH (CH3)3COH(aq) + Br (aq) occurs in three elementary steps; determine the rate law for the + - + + + - equation: 1. (CH3)CBr (CH3)3C + Br is slow 2. (CH3)3C + H2O (CH3)3COH2 is fast 3. (CH3)3COH2 + OH (CH3)3COH(aq) + H2O is fast - - A. Rate = k[(CH3)CBr] B. Rate = k[OH ] C. Rate = k[(CH3)CBr] [OH ] D. Rate = k[(CH3)CBr]² 46

Chemistry Notes

Answer: A + - Reasoning: The first equation,(CH3)CBr (CH3)3C + Br , is the limiting reaction. The rate equation is based on the reactants only. (CH3)CBr is the only reactant.

81. Which statement about equilibrium is not true? A. Equilibrium shifts to minimize the impact of changes. B. Forward and reverse reactions have equal rates at equilibrium. C. A closed container of air and water is at a vapor-liquid equilibrium if the humidity is constant. D. The equilibrium between solid and dissolved forms is maintained when salt is added to an unsaturated solution.

Answer:D Reasoning: There is no solid form of the salt in an unsaturated solution.

82. Which statement(s) about reaction rates are true? I. Catalysts shift equilibrium to favor product formation. II. Catalysts increase the rate of forward and reverse reactions. III. A greater temperature increases the chance that a molecular collision will overcome a reaction’s activation energy. IV. A catalytic converter contains a homogeneous catalyst. A. I and II B. II and III C. II, III, and IV D. I, III and IV

Answer: B Reasoning: Statement IV is not true. A catalytic converter is a heterogeneous catalyst (the catalyst is solid the reaction is gaseous and liquid), so therefore C and D are not correct. Catalysts do not favor the forward or reverse reactions; they lower the activation energy required, so statement I is untrue.

83. Write the equilibrium expression Keq for the reaction: CO2(g) + H2(g) ↔ CO(g) + H2O(l).

Answer: D Reasoning: In an expression for the equilibrium constant, reactants go on the bottom, products on the top (eliminating choice B). The stoichiometric coefficients for each reactant/product are the powers. In the case of this equation, all the coefficients are 1. In the case of a heterogeneous reaction, (such as this, where all are gas except water which is a liquid) the concentrations of pure liquids or solids (as opposed to aqueous) are not included.

84. What could cause this change in the energy diagram of a reaction?

A. Adding a catalyst to an endothermic reaction B. Removing a catalyst from an endothermic reaction C. Adding a catalyst to an exothermic reaction D. Removing a catalyst from an exothermic reaction

Answer:B Reasoning: The reaction shown is endothermic (absorbing heat energy). The activation energy increased, so therefore the catalyst was removed.

-10 85. BaSO4 (Ksp = 1x10 ) is added to pure H2O. How much is dissolved in 1 L of saturated solution? A. 2 mg B. 10 g C. 2 g D. 100 pg

Answer: A

Chemistry Notes

2+ 2- Reasoning: Ksp = [cation][anion]; in this case the cation is Ba and the anion is SO4 . The concentration the barium ion and the 2 -5 sulfate ion is 1:1 so: Ksp = [cation] ; or the concentration = 1 x 10 M. The molar mass of barium sulfate is 239 g/mole so 0.00239 g or 2.39 mg.

86. The exothermic reaction 2NO(g) + Br2(g) ↔ 2NOBr(g) is at equilibrium. According to LeChatelier’s Principle: A. Adding Br2 will increase [NO]. B. An increase in container volume (with temperature constant) will increase [NOBr]. C. An increase in pressure (with temperature constant) will increase [NOBr]. D. An increase in temperature (with constant pressure) will increase [NOBr].

Answer: C Reasoning: Adding Br2 will favor the increased production of NOBr. Increasing pressure pushes the reaction away from the side with more moles of gas; this means that increasing volume (decreasing pressure) will favor the reverse reaction (less NOBr) and increasing pressure will favor the forward reaction (more NOBr). Increasing temperature favors exothermic forward reactions.

3 87. At a certain temperature, T, the equilibrium constant for the reaction 2NO(g) ↔ N2(g) + O2(g) is Keq = 2 x 10 . If a 1.0 L container at this temperature contains 90 mM of N2, 20 mM of O2 and 5 mM NO, what will occur? A. The reaction will make more N2 and O2. B. The reaction is at equilibrium C. The reaction will make more NO D. The temperature, T, is required to solve this problem.

Answer: A 2 2 Reasoning: Keq = [N2][O2] / [NO] for this equation. At this point: Keq = (90)(20) / 5 = 180/25 < 2000; so the reaction is not at equilibrium. Currently, there is too little reactants for the amount of product, so the reaction will favor the reverse reaction.

88. Which statement about acids and bases is not true? A. All strong acids ionize in water. B. All Lewis acids accept an electron pair C. All Bronsted bases use OH- as a proton acceptor D. All Arrhenius acids form H+ ions in water.

Answer: C

89. Which of the following are listed from weakest to strongest acid? - 2- A. H2SO3, H2SeO3, H2TeO3 B. HBrO, HBrO2, HBrO3, HBrO4 C. HI, HBr, HCl, HF D. H3PO4, H2PO4 , HPO4

Answer: B Reasoning: Series D shows a polyprotic series; strength in a polyprotic series decreases with the most protons. Series C shows a series of strong acids except for hydrofluoric acid (hydrogen fluoride) which is a weak acid. Series A shows a series with the same available oxygen and hydrogen but the central atom is decreasing in Electronegativity so the series is getting weaker. Series B shows the same central atom with increasing oxidation number, so increasing strength of acidity.

+ 90. NH4F is dissolved in water. Which of the following are conjugate acid/base pairs present in the solution? I. NH4 /NH4OH. II. - + - HF/F . III. H3O /H2O. IV. H2O/OH . A. I, II, and III B. I, III, and IV C. II and IV D. II, III and IV

Answer: D + - Reasoning: The NH4F dissolves completely in water into NH4 and F . The conjugate acid/base pair must have the form HX/X where X has one charge lower than HX. HF/F-: The F- ion is a weak base and HF is a weak acid (The fluorine ion accepts a proton). + + + NH4 /NH3: NH4 is a weak acid and NH3 as its conjugate base (NH4 donates a proton and becomes NH3) – therefore statement I. is + - untrue. H3O /H2O and H2O/OH : The water acts as a base AND acid.

+ - + - + - - + Complete equations: NH4F(s) + H2O(l) ↔ NH4 (aq) + F (aq) + H (aq) + OH (aq) : NH4 + OH ↔ NH3(aq) + H2O(l) : F + H3O ↔ HF(aq) + - + + H2O(l) : 2H2O ↔ H3O (aq) + OH (aq)

91. What are the pH and pOH of 0.010M HNO3(aq)? A. pH=1.0, pOH=9.0 B. pH=2.0, pOH=12.0 C. pH=2.0, pOH=8.0 D. pH=8.0, pOH=6.0 48

Chemistry Notes

Answer: B Reasoning: Convert concentration into scientific notation: 0.010 = 1.0 x 10-2. pH = - log[x] = -log(1E-2) = -(-2) = 2. The sum of pH + pOH = 14.0 for solutions in equilibrium. pOH = 14 – 2 = 12.

Solving Logarithms (Base 10) without a calculator: Log base 10 of 1000 or log(1000) means 10y =1000 or 10x10x10 = 1000, so y = 3, log(1000) = 3 Log base 10 of 1 x 10-2 or log(10-2) means 10y = 10-2, so y = -2

92. What is the pH of a buffer made of 0.128M sodium formate (HCOONa) and 0.072M formic acid (HCOOH)? The pKa of formic acid is 3.75. A. 2.0 B. 3.0 C. 4.0 D. 5.0

Answer: C pKa -3.75 3.75 -4 + Reasoning: pKa = -log(Ka), so Ka = 10 = 10 = 1/(10 ) = 1.78 x 10 . The dissociation equation for the acid is HCOOH → H + - + - -4 COOH ; so the formula for the equilibrium constant will be Ka = [H ][COOH ] / [HCOOH] = 1.78 x 10 . In this equation, the concentration of anions will be equal to the concentration of buffer base or 0.128M. Solving for hydrogen ion concentration: [H+] = - -4 -4 -4 Ka[HCOOH] / [COOH ] = (1. 78 x 10 )(0.072) / (0.128) = 1.0 x 10 M. So pH = -log(1.0 x 10 ) = 4.

Solving without a calculator: Estimate the order of magnitude of the value of the exponent where 1/(103.75) ≈ 1/(104) ≈ 1/10000 = 1 x 10-4 (Remember that the true value is slightly GREATER than 1 x 10-4) + - -4 -4 Determine the approximate value of pH [H ] = Ka[HCOOH] / [COOH ] = (1 x 10 )(.072) / (0.128) = (1 x 10 )(.6) = 0.0006 This value is a slight underestimate; so the answer is in the range approaching 10-4. The pH of 10-4 is 4; go with choice C.

93. A sample of 50.0 mL KOH is titrated with 0.100M HClO4. The initial buret reading is 1.6 mL and the reading at the endpoint is 22.4 mL. What is the concentration of KOH? A. 0.0416M B. 0.0481M C. 0.0832M D. 0.0962mM

Answer: A Reasoning: Determine the volume of the chloric acid 22.4-1.6 = 20.8 mL. Set up the titration equation where 50x = 0.1(20.8). Solve for x = 0.041M (approx).

94. Rank the following from lowest to highest pH. Assume a small volume for the component given in moles. I. 0.01 mol HCl added to 1L water. II. 0.01 mol HI added to 1L of an acetic acid/sodium acetate solution at a pH of 4.0. III. 0.01 mol ammonia added to 1 L of water. IV. 0.1 mol nitric acid (HNO3) added to 1 L of a 0.1M calcium hydroxide (Ca(OH)2)solution. A. I

Answer: A Reasoning: The greatest pH has the least concentration of free hydrogen ions released during the dissolution process. I. 0.01 mol HCl in 1 L is a molarity of 0.01M, so the concentration of hydrogen ions is 0.01M (pH = 2). II. 0.01 mol HI added to 1 L to water would also have a pH of 2; but since the solution is buffered, the pH will be between 2 and 4. III. 0.01 mol ammonia in 1 L of water has a concentration of 0.01M but this is an hydroxide concentration, so the pOH approximately 2, and the pH approximately 12. (ammonia is a weak base so it won’t full dissolve) IV. The last choice is a neutralization reaction: 2HNO3 + Ca(OH)2 → Ca(NO3) 2 + H2O. The concentrations are equal but 2 moles of the acid reacts with 1 mole of the base, leaving a base concentration of 0.05 – more basic than choice III.

95. The curve below resulted from the titration of a ______with a ______titrant. A. weak acid:strong base B. weak base:strong acid C. strong acid:strong base D. strong base:strong acid

Answer: D Reasoning: The solution is strongly basic initially (pH>7) and strongly acidic at the end. 49

Chemistry Notes

96. Which statement about thermochemistry is true? A. Particles in a system move about less freely at high entropy. B. Water at 100°C has the same internal energy as water vapor at 100°C. C. A decrease in the order of a system corresponds to an increase in entropy. D. At its sublimation temperature, dry ice has a higher entropy than gaseous carbon dioxide.

Answer: C Reasoning: Entropy increases as disorder increases, the lack of molecular bonds means steam has more energy than boiling water, gaseous states always have more entropy than solid states.

97. What is the heat change of 36.0 g water at atmospheric pressure when its temperature is reduced from 125°C to 40°C? Use the following data: heat capacity of solid water 37.6 J•mol-1•°C-1 heat capacity of liquid water 75.3 J•mol-1•°C-1 heat capacity of water vapor 33.1 J•mol-1•°C-1 heat of fusion 6.02 kJ•mol-1 heat of vaporization 40.67 kJ•mol-1 A. -92.0 kJ B. -10.8 kJ C. 10.8 kJ D. 92.0 kJ

Answer: A Reasoning: Reducing temperature is a reduction in heat. The only logical answers are A or B. 36 g of water at 18 g/mol is 2 moles. So calculating heat = nCΔT + nHv + nCΔT = n(CΔT + Hv + CΔT) = 2[(33.1)(25) + 40670 + (75.3)(60)] = 2[827.5 + 40670 + 4518.0] or approximately = 2[1000 + 40000 + 5000] = 2[46000] = 92000 J lost or -92 kJ.

98. What is the standard heat of combustion of methane gas? Use the following data:

Standard heat of formation: CH4(g) -74.8 kJ/mol Standard heat of formation: CO2(g) -393.5 kJ/mol Standard heat of formation: H2O(g) -285.8 kJ/mol A. -890.3 kJ/mol B. -604.6 kJ/mol C. -252.9 kJ/mol D. -182.5 kJ/mol

Answer: B Reasoning: The heat of formation of an element in its most stable for (in this example the diatomic oxygen) has a heat of formation of 0 kJ/mol. Establish and balance the combustion equation: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g). The heat of combustion is equal to the ΔHreaction = Hproducts - Hreactants = 2(ΔHwater) + (ΔHcarbon dioxide) - 2(ΔHoxygen) - (ΔHmethane) = (2)(-285.8) + (-392.5) – (2)(0) – (-74.8) = - 890.3 kJ.

99. Which reaction creates products at a lower total entropy than the reactants? + - A. Dissolution of table salt: NaCl(s) → Na (aq) + Cl (aq) B. Oxidation of iron: 4Fe(s) + 3O2 → 2FeO3 C. Dissociation of ozone: O3(g) → O2(g) + O(g) D. Vaporization of butane: C4H10(l) → C4H10(g)

Answer: B Reasoning: In the dissolution, dissociation, and vaporization, the products are less organized and have more energy. The oxidation reaction decreases entropy.

100. Which statement about reactions is true? A. All spontaneous reactions are both exothermic and cause an increase in entropy. B. An endothermic reaction that increases the order of the system cannot be spontaneous. C. A reaction can be non-spontaneous in one direction and also non-spontaneous in the opposite direction. D. Melting snow is an exothermic process.

Answer: B Reasoning: Spontaneous reactions that increase entropy can be endothermic (Gibb’s Free Energy) or exothermic. Spontaneous reactions that decrease entropy are always exothermic. The only 50

Chemistry Notes combination that cannot be spontaneous is an endothermic reaction that decreases entropy.

101. 10. kJ of heat are added to one kilogram of iron at 10.°C. What is its final temperature? The specific heat of iron is 0.45 J/g°C. A. 22°C B. 27°C C. 32°C D. 37°C

Answer: C Reasoning: From the answer choices, it is evident that it is unnecessary to consider phase changes. ΔT = Q/mC = (10,000 J) / (1000 g x 0.45 J/g °C) = 22°C; T = Ti + ΔT = 10°C + 22°C = 32°C

102. Which reaction is not a redox process? A. Combustion of octane: 2C8H18 + 25O2 → 16CO2 + 18H2O B. Depletion of a lithium battery: Li + MnO2 → LiMnO2 C. Corrosion of aluminum by acid: 2Al + 6HCl → 2AlCl3 + 3H2 D. Taking an antacid for heartburn: CaCO3 + 2HCl → CaCl2 + H2CO3 → CaCl2 + CO2 + H2O

Answer: D Reasoning: In a redox reaction, the oxidation of one element increases (oxidizes; more positive) and one decreases (reduces: more (+4) (-1) (0) (+4) (-2) (+1) (-2) negative). Combustion: 2C8 H18 + 25O2 → 16C O2 + 18H2 O ; the hydrogen oxidizes and the oxygen reduces. Depletion: (0) (+4) (-2) (+1) (+3) (-2) (0) (+1) (-1) (+3) (-1) Li + Mn O2 → Li Mn O2 ; the lithium oxidizes and the manganese reduces. Corrosion: 2Al + 6H Cl → 2Al Cl3 + (0) (+2) (+4) (-2) (+1) (-1) (+2) (-2) (+1) (+4) (-2) 3H2 ; the aluminum oxidizes and the hydrogen reduces. Antacid: Ca (C O3 ) + 2H Cl → Ca Cl2 + H2 C O3 → (+2) (-1) (+4) (-2) (+1) (-2) Ca Cl2 + C O2 + H2 O ; none of the substances changes oxidation states.

103. Given the following heats of reaction: ΔH = -0.3 kJ/mol for Fe(s) + CO2(g)  FeO(s) + CO(g); ΔH = 5.7 kJ/mol for 2Fe(s) + 3CO2(g)  Fe2O3(s) + 3CO(g); and ΔH = 4.5 kJ/mol for 3FeO(s) + CO2(g)  Fe3O4(s) + CO(g). Use Hess’s Law to determine the heat of reaction for: 3FeO3(s) + CO(g)  2Fe3O4(s) + CO2(g). A. -10.8 kJ/mol B. -9.9 kJ/mol C. -9.0 kJ/mol D. -8.1 kJ/mol

Answer: B Reasoning: ΔH = -0.3 kJ/mol for Fe(s) + CO2(g)  FeO(s) + CO(g) ΔH = 5.7 kJ/mol for 2Fe(s) + 3CO2(g)  Fe2O3(s) + 3CO(g) ΔH = 4.5 kJ/mol for 3FeO(s) + CO2(g)  Fe3O4(s) + CO(g) ΔH = X kJ/mol for 3Fe2O3(s) + CO(g)  2Fe3O4(s) + CO2(g)

Flip the second equation to put the Iron (III) oxide on the product side of the equation. (Change sign of heat).

ΔH = -0.3 kJ/mol for Fe(s) + CO2(g)  FeO(s) + CO(g) ΔH = -5.7 kJ/mol for Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g) ΔH = 4.5 kJ/mol for 3FeO(s) + CO2(g)  Fe3O4(s) + CO(g) ΔH = X kJ/mol for 3Fe2O3(s) + CO(g)  2Fe3O4(s) + CO2(g)

Multiply the 2nd and 3rd equations in order to balance with the unknown equation.

ΔH = -0.3 kJ/mol for Fe(s) + CO2(g)  FeO(s) + CO(g) [ΔH = -5.7 kJ/mol for Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g) ]x3 [ΔH = 4.5 kJ/mol for 3FeO(s) + CO2(g)  Fe3O4(s) + CO(g) ]x2 ΔH = X kJ/mol for 3Fe2O3(s) + CO(g)  2Fe3O4(s) + CO2(g)

Multiply the 1st equation in order to balance with the other equations to cancel.

[ΔH = -0.3 kJ/mol for Fe(s) + CO2(g)  FeO(s) + CO(g)]x6 ΔH =-17.1 kJ/mol for 3Fe2O3(s) + 9CO(g)  6Fe(s) + 9CO2(g) ΔH = 9.0 kJ/mol for 6FeO(s) + 2CO2(g)  2Fe3O4(s) + 2CO(g) ΔH = X kJ/mol for 3Fe2O3(s) + CO(g)  2Fe3O4(s) + CO2(g) 51

Chemistry Notes

Check for cancellation across

[ΔH = -1.8 kJ/mol for 6Fe(s) + 6CO2(g)  6FeO(s) + 6CO(g) ΔH =-17.1 kJ/mol for 3Fe2O3(s) + 9CO(g)  6Fe(s) + 9CO2(g) ΔH = 9.0 kJ/mol for 6FeO(s) + 2CO2(g)  2Fe3O4(s) + 2CO(g) ΔH = X kJ/mol for 3Fe2O3(s) + CO(g)  2Fe3O4(s) + CO2(g)

Since the equations now sum to the desired equation, add the heats: -1.8 + -17.1 + 9 = -9.9 kJ/mol.

104. What is the oxidant in the reaction: 2H2S + SO2  3S + 2H2O? A. H2S B. SO2 C. S D.H2O

Answer: B (+1) (-2) (+4) (-2) (0) (+1) (-2) Reasoning: The oxidant is the agent that is reduced. 2H2 S + S O2  3S + 2H2 O ; The oxidation number of both the hydrogen and the oxygen do not change. The sulfur in the hydrosulfuric acid is oxidized to sulfur by the sulfur dioxide; the sulfur dioxide is reduced to sulfur by the hydrosulfuric acid. The oxidant then is the sulfur dioxide.

105. Molten NaCl is subjected to electrolysis. What reaction takes place at the cathode. - - - - + - + - A. 2Cl (l)  Cl2(g) + 2e B. Cl2(g) + 2e  2Cl (l) C. Na (l) + e  Na(l) D. Na (l)  Na(l) + e

Answer: C Reasoning: Removing an electron from the positively charge sodium cation would not leave neutrally charged sodium (eliminating D). In an electrolytic cell, the reduction occurs at the cathode and the ion gains an electron. Since reduction causes the oxidation (ionization) to become more negative. Additional, electrolysis of MOLTEN NaCl starts in the liquid phase.

106. What is the purpose of the salt bridge in an electrochemical cell? A. To receive electrons from the oxidation half-reaction B. To relieve the buildup of positive charge in the anode half-cell. C. To conduct electron flow. D. To permit positive ions to flow from the cathode half-cell to the anode half-cell.

Answer: D Reasoning: The anode receives electrons from the oxidation half-reaction and the circuit conducts the electron flow to the cathode which is then supplied to the reduction half-reaction.

107.Given: Eº = -2.37V for Mg2+(aq) + 2e-  Mg(s) and Eº = 0.80V for Ag+(aq) + e-  Ag(s); what is the standard potential of a voltaic cell composed of a piece of magnesium dipped in 1M Ag+ solution and a piece of silver dipped in 1M Mg2+ solution? A. 0.77V B. 1.57V C. 3.17V D. 3.97V

Answer: C Reasoning: The terminal with the more positive voltage will have the reduction reaction and the more negative will have the oxidation. E°cell = E°cathode - E°anode = 2.37V + 0.80V = 3.17V

108. A proper name for this hydrocarbon is: A. 4,5 dimethyl-6-hexene B. 2,3 dimethyl-1-hexene C. 4,5 dimethyl-6-hexyne D. 2-methyl-3-propyl-1-butene

Answer: B Reasoning: Determining the longest straight chain, there are 6 (hex) and 1 double bond, so hexene is the base of the molecule. There are 2 attachments of CH3 (methyl) on the hexene. Using the lowest numbering, the double bond is on the 1st carbon, the first methyl is on the 2nd carbon and the second methyl is on the third carbon: 2,3 dimethyl 1-hexene.

Chemistry Notes

109. An IUPAC approved name for this molecule is: A. Butanal B. Propanal C. Butanoic Acid D. Propanoic Acid

Answer: C Reasoning: There are four carbons in the straight chain, so this is a but- molecule. On one end is an –OH and on another carbon there is a double bond to an oxygen. This family is the carboxyl acids. So the molecule is named Butanoic Acid. The other two names are for aldehydes which would only have the double oxygen bond but no hydroxide.

110. Which molecule has a systematic name of methyl ethanoate?

Answer: A Reasoning: -oates are esters; esters have a oxygen between two chains and a second oxygen on a double bond on the –oate chain. This eliminates C; D is also eliminated since the name would have to specific two double bonds. A and B are both esters. A is methyl ethanoate; B would be ethyl methanoate.

111. This compound contains an: A. alkene, carboxylic acid, ester, and ketone B. aldehyde, alkyne, ester, and ketone C. aldehyde, alkene, carboxylic acid, and ester D. acid anhydride, aldehyde, alkene, and amine

Answer: C Reasoning: The benzene shown has an attachment of a carboxylic acid. In addition, benzenes are an alkene molecule (double carbon bonds). The end chain with the “tail” of the double bonded oxygen is an aldehyde. The remaining double bond to the oxygen on the benzene would be a ketone BUT there is another single bond to oxygen (which Carboxylic Acid would be an ether) so the combination is an ester.

112. Which group of scientists made contributions in the same area of chemistry? Aldehyde A. Volta, Kekule, Faraday, London B. Hess, Joule, Kelvin, Gibbs C. Boyle, Charles, Arrhenius, Pauli D. Davy, Mendeleev, Ramsay, Galvani

Answer: B

113. Which of the following pairs are isomers?

II. Pentanal 2-pentanone

A. I and IV B. II and III C. I, II and III D. I, II, III and IV

Answer: B Reasoning: I is NOT a cis-trans isomer since there is no double bond between the nitrogens and they can freely rotate. Pentanal is a 5-carbon chain with a double bond oxygen atom on one end; 2 pentanone is a double bond oxygen atom in the middle. IV is just a rotated molecule.

Chemistry Notes

114. Which instrument would be most useful for separating two different proteins from a mixture. A. UV/Vis spectrophotometer B. Mass spectrometer C. Gas Chromotograph D. Liquid Chromotograph

Answer: D Reasoning: Spectrophotometer measures the visible light that can pass through – the proteins will not be very distinct. The mass spectrometer will measure the molecular weights but will not separate them. The proteins are not gases.

115. Classify these biochemicals.

A. I-nucleotide, II-sugar, III-peptide, IV-fat B. I-disaccharide, II-sugar, III-fatty acid, IV-polypeptide C. I-disaccharide, II-amino acid, III-fatty acid, IV-polysaccharide D. I-nucleotide, II-sugar, III-triglyceride, IV-DNA

Answer: A Reasoning: I. Phosphate, sugar and amine = nucleotide, II. CnH2nOn is a sugar. III. is a peptide. IV is a fatty acid.

116. You create a solution nof 2.00 g/mL of a pigment and divide the solution into 12 samples. You give four samples each to three teams of students. They use a spectrophotometer to determine the pigment concentration. Here is their data. Which of the following is true?

Team Concentration ( g/mL) Sample 1 Sample 2 Sample 3 Sample 4 1 1.98 1.93 1.92 1.88 2 1.70 1.72 1.69 1.70 3 1.78 1.99 2.87 2.20 A. Team 1 has the most precise data. B. Team 3 has the most accurate data in spite of it having low precision. C. The data from team 2 is characteristic of a systematic error. D. The data from team 1 is more characteristic of random error than the data from team 3.

Answer: C Reasoning: Team 2 has the greatest precision and Team 1 has the greatest accuracy. Team 2 is systematically low; Team 3’s data is all over the place (random error).

117. Which pair of measurements have an identical meaning? A. 32 micrograms and 0.032 g B. 26 nm and 2.60x10-8 m C. 3.01x10-5 m3 and 30.1 ml D. 0.0020 L and 20 cm3

Answer: C Reasoning: A is off by a factor of 1000. B is off in terms of sig figs. D is off by factor of 10.

118. Match the instrument with the quantity it measures. I. Eudiometer II. Calorimeter III. Manometer IV. Hygrometer A. I-volume, II-mass, III-radioactivity, IV-humidity B. I-volume, II-heat, III-pressure, IV-humidity C. I-viscosity, II-mass, III-pressure, IV-surface tension D. I-viscosity, II-heat, III-radioactivity, IV- surface tension.

Answer:B 54

Chemistry Notes

Reasoning: A calorimeter measures heat, a manometer measures pressure, a hygrometer measures humidity. Therefore a eudiometer must measure volume.

119. Four nearly identical gems from the same mineral are weighed using different balances. Their masses are: 3.4533 g, 3.459 g, 3.4656 g and 3.464 g. The four gems are then collected and added to a volumetric cylinder containing 10.00 mL of liquid and a new volume of 14.97 mL is read. What is the average mass of the four stones and what is the density of the mineral. A. 3.4650 g, 2.78 g/mL B. 3.460 g, 2.79 g/mL C. 3.4605 g, 2.78 g/mL D. 3.461 g, 2.79 g/mL

Answer:B

120. Which list includes equipment that would not be used in vacuum filtration? A. Rubber tubing, Florence flask, Buchner funnel B. Vacuum pump, Hirsch Funnel, Rubber stopper with a single hole C. Aspirator, Filter paper, filter flask D. Lab Stand, Clamp, filter trap

Answer: A

121. Which of the following statements about lab safety is not true. A. Corrosive chemicals should be stored below eye level. B. A chemical splash on the eye or skin should be rinsed for 15 minutes in cold water. C. MSDS means “material safety data sheet.” D. A student should “stop, drop and roll” if their clothing catches fire in the lab.

Answer: D. Reasoning: Use a safety shower.

122. Which of the following lists consists entirely of chemicals that are considered safe enough to be in a high school lab? A. hydrochloric acid, lauric acid, potassium permanganate, calcium hydroxide B. ethyl ether, nitric acid, sodium benzoate, methanol C. cobalt (II) sulfide, ethylene glycol, benzoyl peroxide, ammonium chloride D. picric acid, hydrofluoric acid, cadmium chloride, carbon disulfide

Answer: A Reaoning: Picric acid, hydrofluoric acid, ethers and benzoyl peroxide are all hazardous.

123. The following procedure was developed to find the specific heat capacity of metals: 1. Place pieces of the metal in an ice-water bath so that their initial temperature is 0ºC. 2. Weigh a Styrofoam cup. 3. Add water at room temperature to the cup and weigh it again. 4. Add a cold metal from the bath to the cup and weigh the cup a third time. 5. Monitor the temperature drop of the water until a final temperature at thermal equilibrium is found. ______is also required as additional information in order to obtain heat capacities for the metals. The best control would be to follow the same protocol expect to use ______in step 4 instead of cold metal. A. The heat capacity of water/a metal at 100 ºC B. The heat of formation of water/ice from the 0ºC bath C. The heat capacity of ice/glass at 0 ºC D. The heat capacity of water/water from the 0ºC

Answer: D Reasoning: It is necessary to know how liquid water responds to heat flow and water from the bath would allow you to see how the homogeneous mixture responds to mixing.

124. Which statement about the impact of chemistry on society is not true? A. Partial hydrogenation creates transfat. B. The Haber Process incorporates nitrogen from the air into molecules for agricultural use. C. The carbon dioxide concentration in the atmosphere has decreased in the last ten years. D. The concentration of ozone-destroying chemicals in the stratosphere has decreased in the last ten years.

Answer: C 55

Chemistry Notes

125. Which statement about everyday applications of chemistry is true? A. Rainwater found near sources of air pollution will most likely be basic. B. Batteries run down more quickly at low temperatures because the chemical reactions are proceeding more slowly. C. Benzyl alcohol is a detergent used in shampoo D. Adding salt decreases the time required for water to boil

Answer: B Reasoning: Rainwater will more likely be acidic. Benzyl Alcohol is not polar enough to be a detergent. Adding salt increases the time for water to boil.

Question Count Category Question #’s Calculations Required Kinetic Molecular Theory and Gas Laws 1, 4, 6, 7, 8, 9, 11, 17, 74 (7%) Gas Laws Properties of Solids, Liquids, Gases 2, 3, 5, 10 (3%) Volume, Number of Molecules Properties of Compounds; intermolecular 12, 13, 14, 15 (3%) bonds Diffusion Rates and Partial Pressures 16, 18, 19, 20, 21, 22 (5%) Graham’s Law, Partial Pressure Thermal Chemistry, triple points, phase 24, 25, 26, 27 (3%) changes, Endothermic and Exothermic Solubility, Molarity, Molality and 28, 29, 30, 31, 32, 33, 34, 35, 36, 37, 38, Molarity, Molality, Normality, Weight Normality, Colligative Properties, Osmotic 39, 40, 41, 42, 43 Percentage, Mole Fractions, Raoult’s Law, Pressure, Raoult’s Law (13%) Boiling Pt Elevation Nuclear Decay, Half-Life, Applications, 44, 45, 46, 47, 48, 49, 50, 53, 54 Half-Life, Average Atomic Mass Average Atomic Mass (7%) Scientists and Theories 51, 52, 56. 112 (3%) Electron Orbitals, Lewis Dot Structures, 55, 57, 63, 64, 65 (4%) Conjugate Molecules Periodic Trends, Electronegativity, Atomic 58, 59, 60, 61, 62, 66 (5%) Radius, Polarity of Molecules, VSPER IUPAC Naming, Chemical Composition, 67, 68, 69, 70, 71, 72, 73, 75, 76, 77, 78 % composition, mass produced, volume Chemical Reactions, Limiting Reagents. (9%) produced Rate Laws, Equilibrium, Catalysts, 79, 80, 81, 82, 83, 84, 85, 86, 87 Rate law, Equilibrium Constant, Equilibrium Expressions, Solution (7%) LaChateilier’s Principle Constants, La Chateilier’s Principle Acids and Bases, Conjugate Acid/Base 88, 89, 90, 91, 92, 93, 94, 95 pH, pH of buffer solutions, titration pairs, Acid Strength, pH and pOH (6%) Thermochemistry, Heat of Combustion, 96, 97, 98, 99, 100, 101, 103 Heat Flow, Heat of Combustion, Heat of Entropy, Gibbs Free Energy (6%) Reaction Reduction, Oxidation Reactions, 102, 104, 105, 106, 107 (4%) Oxidants, Standard Cell Potential Electrochemistry Hydrocarbons, Isomers 108, 109, 110, 111, 113, 115 (5%) Measurements, Equipment, Safey and Lab 23, 114, 116, 117, 118, 119, 120, 121, 122, Procedures 123, 124, 125 (10%)