Chemistry Check Off List

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CORE CURRICULUM CHECK-OFF LIST REGENTS CHEMISTRY

The following information is everything you need to know and be able to do to attain Mastery of the Regents Chemistry Curriculum. The bold, underlined words are important vocabulary words that you should be able to define and use properly in explanations. This is a study guide for what you will be tested on throughout the year. The objectives are divided into categories of “Knowledge” (what you have to know) and “Application” (what you have to be able to do). Check off each objective when you fully understand it. If you do not understand an objective, ask questions before you are tested on it.

I. MATHEMATICS ANALYSIS & GRAPHING o Identify independent and dependent variables in an experiment and correctly plot them on an axis

Example hypothesis: Chemistry students who do their homework will have higher test scores than 1 students who do not do their homework. .  X-axis (horizontal): the independent variable is the one that is manipulated by the experimenter. (“The one I change.” – do homework/not do homework)  Y-axis (vertical): the dependent variable is the one that changes based on the independent variable. (The data you collect – test scores) o Express uncertainty in measurement by properly using significant figures o Identify the number of sig. figs. in data 2 o Round to the correct number of sig. figs. on calculations: .  Addition and Subtraction (round to least precise [farthest left] place value)  Multiplication and Division (round to the lowest # of digits in data)  Combined Rules (Add/Subtract first, then Multiply/Divide) 3 o Identify relationships between variables from data tables and graphs (direct or inverse . relationships) 4 o Understand what is meant by conditions of Standard Temperature and Pressure (STP) . (Table A) o Recognize and convert between units on various scales of measurement (Tables C, D, and T)

 Temperature: Celsius ↔ Kelvin  Length: meters ↔ centimeters ↔ 5 millimeters .  Mass: grams ↔ kilograms  Pressure: kilopascals ↔ atmospheres  Thermal Energy: joules ↔ kilojoules  Amount of Substance: GFM = 1 mole = 6.02x1023 particles o grams ↔ moles ↔ atoms or molecules 6 o Use the density equation on Table T to solve for density, mass, or volume, given the other two . values 7 o Calculate percent error (Table T) . II. ATOMIC CONCEPTS Knowledge Application o Explain what happened during the gold- foil experiment and what it showed. o Atoms are the basic unit (building block) of Gold foil was bombarded (hit) with matter. positively charged alpha particles. Most alpha particles passed through the gold o Atoms of the same kind are called elements. 1. foil, but some were deflected. This showed The modern model of the atom has developed o that: over a long period of time through the work of 1. the atom is mostly empty space many scientists. 2. the nucleus is small, positively charged, and located in the center of the atom o The three subatomic particles that make up an atom are protons, neutrons, and electrons. The proton is positively charged, the neutron has no charge, and the electron is negatively 2. o charged. (This is also referenced on Table O. Check it out!) o Each atom has a nucleus with an overall positive charge, made up of protons and o Determine the nuclear charge of an atom. 3. neutrons. (equal to the number of protons in the o The nucleus is surrounded by negatively nucleus) charged electrons. o Atoms are electrically neutral, which means Determine the number of protons or that they have no charge (# protons = # o electrons) electrons in an atom or ion when given one of these values. (Periodic Table) o Ions are atoms that have either lost or gained Compare the atomic radius and ionic electrons and are either positively or negatively o 4. charged radius of any given element Ex: A chloride ion has a larger radius o When an atom gains one or more electrons, it than a chlorine atom because the ion has becomes a negative ion and its radius an extra electron. A sodium ion has a increases. smaller radius than a sodium atom When an atom loses one or more electrons, it o because the ion has lost an electron. becomes a positive ion and its radius decreases. o The mass of each proton and each neutron is o Calculate the mass of an atom given the approximately equal to one atomic mass unit number of protons and neutrons 5. (AMU). o Calculate the number of neutrons or o An electron is much less massive (has almost protons, given the other value no mass) compared to a proton or neutron. o In the modern model of the atom, the WAVE-MECHANICAL MODEL (electron cloud), the 6. electrons are in orbitals (clouds), which are defined as regions of most probable electron location. o Each electron in an atom has a specific amount of energy. 7. o Electrons closest to the nucleus have the lowest energy. As an electron moves away from the nucleus, it has higher energy. 8. o Electron configurations show how many o Distinguish between ground state and electrons are in each orbital. excited state electron configurations. o When all of an atom’s electrons are in the (Be careful to keep the same number of orbitals closest to the nucleus, the electrons are electrons when writing the excited state.) in their lowest possible energy states. This is called the ground state. o When an electron in an atom gains a specific amount of energy, the electron is at a higher energy state (excited state) II. ATOMIC CONCEPTS (CONTINUED) Knowledge Application o When an electron returns from a higher energy (excited) state to a lower energy (ground) state, a specific amount of energy is emitted. This emitted energy can be used to identify an o Identify an element by comparing its 9. element. bright-line spectrum to given spectra o The flame test is an example of the bright-line spectrum visible to the naked eye. The color can determine the identity of a positive ion in a compound. o Draw a Lewis electron-dot structure of an o The outermost electrons in an atom are called atom. 10 the valence electrons. In general, the number o Distinguish between valence and non- . of valence electrons affects the chemical valence electrons, given an electron properties of an element. configuration o Atoms of an element that contain the same o Calculate the number of neutrons in an 11 number of protons but a different number of isotope of an element given the isotope’s . neutrons are called isotopes of that element. mass o Given an atomic mass, determine the most abundant o The average atomic mass of an element is the 12 isotope weighted average of the masses of its naturally . o Calculate the atomic mass of an element, occurring isotopes. given the masses and abundance of naturally occurring isotopes

III. THE PERIODIC TABLE Knowledge Application o The placement or location of an element on the Periodic o Explain the placement of an Table gives an indication of physical and chemical unknown element in the 1. properties of that element. periodic table based on its o The elements on the Periodic Table are arranged in order properties of increasing atomic number. 2. o The number of protons in an atom (atomic number) o Interpret and write isotopic identifies the element. This goes at the bottom left corner notations. of the symbol for that element. Ex: o The sum of the protons and neutrons in an atom (mass C-12, C-13, and C-14 are number) identifies an isotope. The mass number is isotopes of the element carbon placed at the top left corner of the symbol for an element OR is placed after the element symbol or name and a dash. Ex: Three different ways to write carbon with a mass number of 14: C, carbon-14, or C-14 o Identify the properties of metals, metalloids, nonmetals o Elements are classified by their properties, and located on and noble gases 3. the periodic table as metals, metalloids, nonmetals, or o Classify elements as metals, noble gases. metalloids, nonmetals, or noble gases by their properties o III. THE PERIODIC TABLE (CONTINUED) Knowledge Application o An element’s atomic radius, first ionization energy, and electronegativity determine its 4. physical and chemical properties o Substances can be differentiated by their physical o Identify and give examples of properties. physical properties 5. o Physical properties of substances include melting point, o Describe the states of the boiling point, density, conductivity, malleability, elements at STP (solid, liquid, solubility, and hardness. or gas). (Table S) o Substances can be differentiated by chemical o Identify and give examples of properties. chemical properties 6. o Chemical properties describe how an element behaves o Describe the difference between during a chemical reaction and include reactivity, physical and chemical flammability, and toxicity. properties of substances o Some elements exist as two or more forms in the same phase. These forms differ in their molecular or crystal structure and therefore in their properties. The word to describe this phenomenon is ALLOTROPE. 7. o Ozone and oxygen gases are allotropes of each other. Ozone is O3 and it is very dangerous to our health. Oxygen gas is O2 and we need it to survive. o Diamonds and graphite (better known as pencil lead) are both forms of the element carbon. They have different molecular structures and very different properties. o Determine the group of an element, given the chemical o For Groups (also called families) 1, 2, and 13-18 on the formula of a compound Periodic Table, elements within the same group have the Ex: A compound has the same number of valence electrons (helium is the formula XCl , element X is in 8. exception) and therefore similar chemical properties. 2 Group 2 o Elements in the same Period (row) have the same o Determine the number of number of principal energy levels (shells) which contain energy levels containing electrons. electrons given an element’s Period and vice versa o The succession of elements within the same GROUP (top o Compare and contrast to bottom) demonstrates characteristic trends: properties of elements within a 9. differences in atomic radius, ionic radius, group or a period for groups 1, electronegativity, first ionization energy, and 2, and 13-18 on the periodic metallic/nonmetallic properties. table 10 o The succession of elements across the same PERIOD o Understand and be able to . (left to right) demonstrates characteristic trends: explain the trends in terms of differences in atomic radius, ionic radius, nuclear charge and electron electronegativity, first ionization energy, and shielding metallic/nonmetallic properties.

IV. MATTER, ENERGY, & CHANGE Knowledge Application o Matter is anything that has mass and volume (takes up space). o Identify specific examples of 1 o Matter cannot be created nor destroyed, only matter as an element, compound, . transformed. or mixture o Matter is classified as a pure substance (element or compound) or as a mixture of substances. IV. MATTER, ENERGY, & CHANGE (CONTINUED) Knowledge Application o Energy is not matter (does not have volume) o Energy can exist in different forms, such as kinetic, 2 potential, thermal (heat), sound, chemical, electrical, o Distinguish between matter and . and electromagnetic. energy o Energy cannot be created or destroyed, only transformed. o During a physical change, particles of matter are rearranged. Examples of physical changes include freezing, melting, boiling, condensing, dissolving, o Differentiate between physical and crystallizing, and crushing into a powder chemical changes in matter 3 o During a chemical change, NEW substances are o Identify and give examples of . formed with new properties. Examples of chemical physical changes and chemical changes include combustion (burning), rusting, and changes in matter neutralizing an acid or base. o Energy can be absorbed or released during physical and chemical changes. o A pure substance (element or compound) has a uniform composition and constant properties throughout o Draw and interpret particle 4 a given sample, and from sample to sample. diagrams for elements, . o Mixtures are composed of two or more different compounds, and mixtures substances that can be separated by physical means. 5 o Elements are substances that are composed of atoms that have the same atomic number. . Elements cannot be broken down by chemical change. o A compound is a substance composed of two or more o Describe differences in ionic and different elements that are chemically combined in a molecular/covalent compounds fixed proportion. o Identify a compound as ionic or o A compound can be broken down by chemical means, molecular/covalent compound 6 such as during a chemical reaction. given its properties . o Two major categories of compounds are ionic and o Name compounds based on their molecular (covalent) compounds. chemical formulas o A chemical compound can be represented by a specific o Determine the formula of a chemical formula and assigned a name based on the compound given its name IUPAC system. o When different substances (elements or compounds) are mixed together and do NOT chemically react, a mixture is formed. o Interpret particle diagrams as o The amounts of substances in a mixture can vary. Each showing homogeneous or 7 substance in a mixture retains its original properties. heterogeneous mixtures . o The composition of a mixture can vary. If the o Give examples of homogeneous substances are uniformly (evenly) distributed and heterogeneous mixtures throughout the mixture, it is called a homogenous mixture. If the substances are unevenly distributed, it is called a heterogeneous mixture. o Differences in properties such as density, particle size, o Describe the processes of filtration, 8 molecular polarity, boiling point, freezing point, and distillation, and chromatography . solubility allow physical separation of the components and the types of mixtures they are of the mixture. used to separate

V. CHEMICAL BONDING Knowledge Application o Atoms bond with other atoms to gain a stable electron 1 o Determine the noble gas configuration an atom will configuration. . achieve when bonding o Noble gases are already stable and tend to not bond/react. o When a chemical reaction takes place, existing bonds must be broken in order for new bonds 2 (and new compounds) to be formed. . o When a bond is broken, energy is absorbed. When a bond is formed, energy is released. o Electron-dot diagrams (Lewis 3 structures) are used to represent o Draw Lewis dot structures for any given element, ion, . the valence electron arrangement or compound in elements, ions, and compounds. o Chemical bonds are formed when o Identify the type of bonding in a compound (ionic or valence electrons are: covalent), given the elements that make it up  transferred from one atom to o Demonstrate bonding concepts using Lewis Dot another (ionic) Structures (electron-dot diagrams) for ionic and  shared between atoms covalent compounds (covalent) Ionic compounds – after the transfer of electrons, the  mobile within a metal positive ion should have no dots. The negative ion should 4 (metallic) have 8 dots around it. Put brackets around the ions. . o Metals tend to react with Check the periodic table to find the charge associated with nonmetals to form ionic bonds. the ion and place this charge outside of the brackets. *Be o Nonmetals tend to react with other sure to include coefficients if there are more than one of nonmetals to form molecular the same kind of ion.* (covalent) bonds. Covalent compounds – each atom in a covalent compound o Ionic compounds containing must end up with 8 dots around it – except for hydrogen polyatomic ions have both ionic (only 2 dots). *If you use lines, remember that one line AND covalent bonds. represents 2 electrons being shared.* o In a multiple covalent bond, more 5 o Draw electron-dot diagrams and give examples of than one pair of electrons is shared . molecules with multiple covalent bonds between two atoms. o Electronegativity indicates how o Distinguish between and give examples of nonpolar strongly an atom of an element covalent bonds and polar covalent bonds attracts electrons in a bond. 6 If two atoms of the same element share electrons, the o The electronegativity difference . bond is nonpolar (ex: H–H) between two bonded determines If atoms of two different elements share electrons, the the degree of polarity in the bond. bond is polar (ex: H–Cl)

o Molecular polarity is determined by the distribution of electrons in a o Distinguish between bond polarity and molecular covalent compound. polarity o Symmetrical distribution of o Draw Lewis Dot Structures for all of the compounds 7 electrons results in (nonpolar) listed to the left, including the diatomic elements . molecules (Ex: CO2, CH4 and all o Determine whether a molecule is polar or nonpolar, diatomic elements) given its structure o Asymmetrical distribution of electrons results in (polar) *SNAP* molecules (Ex: HCl, NH3, H2O)

VI. CHEMICAL FORMULAS, REACTIONS & STOICHIOMETRY Knowledge Application o Chemical formulas are used to represent compounds. o The main types of chemical formulas include: empirical, molecular, and structural. o Determine the empirical o An empirical formula is the simplest whole-number ratio of formula from a molecular atoms in a compound. 1 formula o Molecular formulas are chemical formulas that show the . o Draw structural formulas for actual ratio of atoms in a molecule of that compound. covalent (molecular) Structural formulas can also be used to represent covalent o compounds compounds. These use lines to show covalent bonds between atoms and also show the geometrical arrangement of atoms in the compound. o One mole of any substance is equal to 6.02 x 1023 pieces of o Calculate the molar mass that substance. (gram-formula mass) of a o The formula mass of a compound is equal to the sum of the substance atomic masses of its atoms (units are atomic mass units) o Determine the molecular 2 The molar mass (gram-formula mass) of a substance is formula, given the empirical . o equal to the formula mass in grams – hence “gram-formula formula and the molar mass mass”. o Determine the number of o The mass of one mole of any substance is equal to its molar moles of a substance, given mass (gram-formula mass). its mass and vice versa o Calculate the percent composition of any element 3 o The percent composition by mass of each element in a in a given compound . compound can be calculated mathematically. o Calculate the percent composition of water in a given hydrate 4 o Balanced chemical equations show conservation of matter, o Balance equations, given the . energy, and charge. formulas for reactants and products o The coefficients in a balanced equation can be used to o Calculate simple mole-mole determine mole ratios in the reaction. ratios, given balanced equations o Types of chemical reactions include synthesis, o Identify the different types of 5 decomposition, single replacement, and double chemical reactions, given . replacement. their chemical equations

VII. PHYSICAL BEHAVIOR OF MATTER Knowledge Application o Physical properties of substances can be explained in o Predict relative melting and terms of chemical bonds and intermolecular forces. boiling points of compounds o Intermolecular forces created by an unequal distribution given information about its of charge result in varying degrees of attraction between chemical bonds or strength of 1. molecules. Hydrogen bonding is an example of a intermolecular forces strong intermolecular force that occurs in compounds o Predict relative strength of containing H bonded to F, O, or N atoms. intermolecular forces of a o Physical properties include malleability, solubility, compound given its melting and hardness, melting/freezing point, and boiling point. boiling points VII. PHYSICAL BEHAVIOR OF MATTER – SOLIDS, LIQUIDS, AND GASES Knowledge Application o Draw and interpret a particle o The three phases of matter (solids, liquids, and gases) diagram to differentiate among have different properties. solids, liquids, and gases 2. o The structure and arrangement of particles and their o Explain phase change in terms of interactions determine the physical state (s, l, or g) of a the changes in energy and substance at a given temperature and pressure. intermolecular distances o Heat is a transfer of energy (thermal energy) from a body of higher temperature to a body of lower o Distinguish between heat energy temperature. Heat (thermal energy) is associated with and temperature in terms of the random motion of atoms and molecules. molecular motion and amount of 3. (Heat flows from HOT materials to COLD materials) matter o Temperature is a measure of the average kinetic o Convert between degrees Celsius energy of the particles in a sample of matter. and degrees Kelvin (Table T) Temperature is NOT energy – it is a measure of heat energy. o Identify areas of heating and o The concepts of kinetic and potential energy can be cooling curves that show changes used to explain physical processes that include: fusion in kinetic and potential energy, (melting), solidification (freezing), vaporization heat of vaporization, heat of (boiling, evaporation), condensation, sublimation, and fusion, and phase changes 4. deposition. o Calculate the heat involved in a o The kinetic and potential energy changes involved in phase or temperature change of a these physical processes can be illustrated in a heating sample of matter using Tables B curve or cooling curve. & T and/or a given heating or cooling curve 5. o Physical processes like phase changes can be o Distinguish between endothermic exothermic or endothermic. and exothermic phase changes o Entropy is a measure of the randomness or disorder of o Compare the entropy of different 6. a system. A system with greater disorder has more phases of matter entropy. o The concept of an ideal gas is a model to explain o Given a choice of pressure and behavior of gases. temperature conditions, identify 7. o A real gas is most like an ideal gas when the real gas is those under which gases behave at low pressure and high temperature. most ideally and/or least ideally o The Kinetic Molecular Theory (KMT) states that, for an IDEAL gas, all gas particles a. are in random, constant, straight-line motion b. are separated by great distances relative to their size (have negligible volume) 8. c. have no attractive forces between them d. have collisions that may result in a transfer of energy between them, but the total energy of the system remains constant o Explain the gas laws in terms of KMT. o The Kinetic Molecular Theory (KMT) describes the o Solve problems using the 9. relationships of pressure, volume, temperature, velocity, combined gas law (Table T) frequency, and force of collisions among gas molecules. o Recognize and draw graphs showing P vs. T, V vs. T, and P vs. V 10 o Equal volumes of gases at the same temperature and pressure contain an equal number of . particles. VII. PHYSICAL BEHAVIOR OF MATTER – AQUEOUS SOLUTIONS Knowledge Application 11 o Physical processes, such as a compound dissolving in a o Interpret ∆H values for physical . solution, can be exothermic or endothermic. processes given in Table I o Identify the solute and the o A solution is a homogeneous mixture of a solute solvent in a given solution dissolved in a solvent. o Give examples of different types of solutions o The solubility of a solute in a given amount of solvent is dependent on the temperature, the pressure, and the o Predict the effect of temperature, chemical natures of the solute and solvent. pressure, and nature of solvent on General rules: 12 o solubility for a given solute a. solubility of a solid increases as temperature . o Use a solubility curve to increases (direct relationship) distinguish among unsaturated, b. solubility of a gas decreases as temperature saturated, and supersaturated increases (inverse relationship) solutions c. solubility of a gas increases as pressure increases o Calculate the amount of a specific (direct) solute dissolved at different d. “like dissolves like” – polar solvents dissolve temperatures using Table G polar solutes; nonpolar solvents dissolve nonpolar solutes o Use Table F (Solubility Guidelines) to determine a o Many chemical reactions happen in solution. When compound’s solubility different ionic compounds are mixed together in the o Determine if a precipitate will 13 same solution, a double replacement reaction may occur form when ionic compounds are . and a stable precipitate (insoluble/solid compound) mixed in solution may form. o Write and balance chemical equations for double replacement reactions o Calculate solution concentrations in molarity (M), percent by volume, percent by mass, or parts per million (ppm) o The concentration of a solution may be expressed as: o Describe how you would prepare 14 molarity (M), percent by volume (%v/v), percent by a solution from scratch, given the . mass (%m/v), or parts per million (ppm). desired molarity o Describe how you would dilute a solution of known concentration (must use the equation M1V1 = M2V2) o The addition of a nonvolatile solute to a solvent causes the boiling point of the solution to increase and the o Compare the freezing and boiling 15 freezing point of the solution to decrease. points of solutions of different . o The greater the concentration of solute particles, the concentration greater the increase in b.p. and decrease in f.p.

VIII. KINETICS AND EQUILIBRIUM Knowledge Application 1 o The Collision Theory states that a chemical reaction is most likely to occur if reactant particles . collide with the proper energy and orientation. o Use the Collision Theory to explain how factors such as temperature, surface area, o The rate (speed) of a chemical reaction and concentration influence the rate of depends on several factors: temperature, reaction concentration, nature of reactants, surface Ex: Increasing the temperature, surface area, and the presence of a catalyst. area, or concentration all lead to an 2 o Ionic compounds generally react faster than increase in the rate of a reaction because . covalent (molecular) compounds they all increase the number of effective o A catalyst provides an alternate reaction collisions between reactant particles. pathway, which has lower activation energy o Explain, in terms of the number of bonds than an uncatalyzed reaction. broken, why ionic compounds generally react faster than covalent compounds o Explain how a catalyst speeds up a reaction 3 o Energy released or absorbed during a chemical o Read and interpret a potential energy . reaction can be represented by a potential diagram o Draw and label the following parts of a energy diagram. potential energy diagram for both an o The difference in PE of the products and endothermic and exothermic reaction reactants is called the heat of reaction (  H)  PE of reactants and PE of products H = PE products – PE reactants  heat of reaction (H) o H values for many chemical reactions are  activation energy (for both the forward listed in Table I and reverse reactions)  activation energy with a catalyst present o Chemical and physical changes can reach 4 equilibrium o Distinguish between examples of physical . o Saturated solutions are examples of systems equilibria and chemical equilibria in physical equilibria (aq ↔ s) o At equilibrium, the rate of the forward o Describe what is happening to the reaction equals the rate of the reverse concentrations or amounts of reactants and 5 reaction and the measurable quantities of products in a system at equilibrium . reactants and products remain constant at o Describe the rates of opposing reactions in a equilibrium*CARE* system at equilibrium o LeChatelier’s principle can be used to predict the effect of a stress (such as a change in pressure, volume, concentration, or o Describe, in terms of LeChatelier’s temperature) on a system at equilibrium. principle, the effects of stress on a given o According to LeChatelier’s principle, a system system at equilibrium, including: at equilibrium will “shift” to reduce the  Changing the effects of a stress placed on the system. It will temperature/heating/cooling 6 “shift” AWAY from an INCREASE and will  Changing the concentration of a . “shift”toward a decrease in concentration or reactant or product temperature (“shift” means that either the  Changing the pressure or volume (this forward or the reverse reaction will be affects systems involving gases) “favored” (go faster) until the rates are again o Also be able to explain why any shifting equal and equilibrium is re-established). occurs in terms of Collision Theory o Changing the pressure or volume only affects systems that contain gases 7 o Systems in nature tend to undergo changes toward lower energy and higher entropy. . IX. ORGANIC CHEMISTRY Knowledge Application o Organic compounds contain carbon atoms which bond to one another in chains, rings, and 1 networks to form a variety of structures. . o Organic compounds are named using specific IUPAC rules and can be represented using molecular formulas, structural formulas, or condensed structural formulas. 2 o Hydrocarbons are organic compounds that o Draw structural formulas for alkanes, . contain only carbon and hydrogen. *The C–H alkenes, alkynes, given their IUPAC bonds are considered to be nonpolar covalent names bonds.* o Name a hydrocarbon (IUPAC rules), given o Saturated hydrocarbons contain only single its molecular formula, structural formula, or C–C bonds. condensed structural formula o Unsaturated hydrocarbons contain at least o Identify whether a hydrocarbon is saturated one double or triple carbon-carbon bond. or unsaturated, given its IUPAC name, structural formula, condensed structural formula, or general formula (see Table Q)

o In a multiple covalent bond, more than one o Determine the TOTAL number of electrons 3 pair of electrons is shared between two atoms. shared in a covalent bond . In a structural formula, each line represents o Determine the number of electron PAIRS TWO shared electrons (lines) shared in a covalent bond o A functional group is a group of atoms attached to an organic compound that gives distinct physical and chemical properties to o Identify different kinds of functional groups organic compounds having that group attached o Classify an organic compound based on its to it. structural formula, condensed structural o Organic acids, alcohols, esters, aldehydes, formula, or IUPAC name ketones, ethers, halides, amines, amides, o Draw a structural formula with the 4 and amino acids are types of organic functional group(s) on a straight chain . compounds that differ in the type of functional hydrocarbon backbone, given the correct group they have. IUPAC name for the compound o Compounds that have the same functional o Name any of these organic compounds, group have similar physical and chemical given their structural or condensed properties structural formulas (see Table R) Ex: all esters have pleasant odors, all organic acids donate H+ ions in solution, all alcohols have low boiling points, etc. o Determine if two compounds are isomers, o Isomers are organic compounds that have the 5 given their molecular formulas, structural same molecular formula, but different . formulas, condensed structural formulas, or structures and properties. names o Identify types of organic reactions, given o Types of organic reactions include: balanced chemical equations 6 polymerization, substitution, fermentation, o Determine missing reactants or products in . addition, combustion, esterification, and a balanced equation, given the type of saponification. *P.S. FACES* reaction.

X. OXIDATION-REDUCTION REACTIONS Knowledge Application 1 o An oxidation-reduction (redox) reaction o Determine the number of moles of . involves the transfer of electrons from one electrons lost or gained in a redox reaction, species (element or ion) to another. given the other value o The number of electrons lost equals the number of electrons gained (conservation of charge)

o Assign oxidation states to atoms and ions Determine if a reaction is a redox reaction o Oxidation numbers (states) can be assigned to o 2 atoms and ions. given the reaction equation (Hint: any reaction in which an element is alone . o Changes in oxidation numbers indicate that a (uncombined with another element) on one redox reaction has occurred. side, but in a compound on the other side – it’s a redox reaction!) o Losing Electrons is Oxidation (LEO) o Determine which species undergoes 3 o Gaining Electrons is Reduction (GER) reduction (oxidation state goes down) and . o Oxidized and reduced species are ALWAYS on which species undergoes oxidation the LEFT (reactants) side of the equation. (oxidation state goes up) o An oxidation half-reaction shows which o Determine if a given half-reaction is species is oxidized and the number of electrons showing oxidation or reduction it loses (electrons go on the right side of the o Write and balance oxidation and reduction 4 arrow) half-reactions (*remember conservation of . o A reduction half-reaction shows which mass and charge – multiply one or both of species is reduced and the number of electrons the half-reactions to make electrons lost = it gains (electrons go on the left side of the electrons gained if they are not equal at arrow) first) o An electrochemical cell can either be a voltaic cell (a battery) or an electrolytic cell. o Explain, in terms of atoms and ions, why In both voltaic and electrolytic cells the anode loses mass and the cathode gains 5 o  mass . oxidation occurs at the anode (An Ox)  reduction occurs at the cathode (Red Cat) o Compare and contrast voltaic cells with  the anode loses mass electrolytic cells  the cathode gains mass o Given a diagram of a voltaic cell, identify and label the cathode, anode, salt bridge, o A voltaic cell spontaneously converts and the direction of electron flow 6 chemical energy into electrical energy. o Explain the function of the salt bridge and . o The purpose of the salt bridge is to allow for the direction of positive and negative ion the migration of ions between half-cells migration o Write balanced oxidation and reduction half-reactions o An electrolytic cell requires electrical energy to produce a chemical change. Electrolytic o Given a diagram of an electrolytic cell, 7 cells can be used for electrolysis (splitting a identify and label the cathode, anode, and . compound into its elements) and for direction of electron flow. electroplating (coating something with a metal). XI. ACIDS, BASES, AND SALTS Knowledge Application

o The behavior of many acids and bases can be o Know the definitions of Arrhenius acids and explained by the Arrhenius Theory. bases

o Arrhenius acids produce H + (hydrogen o If given the properties, chemical formula, or ions) as the only positive ions in aqueous name, identify a substance as an Arrhenius solution. The hydrogen ion may also be 1 acid or Arrhenius base. (Use Tables K, L, written as H O + and called the hydronium . 3 and T to help you remember these ion. definitions. Arrhenius acids begin with H, Arrhenius bases are metals + hydroxide o Arrhenius bases produce OH – (hydroxide ion(s). *Don’t be fooled by alcohols, which ions) as the only negative ion is aqueous also end in OH, but contain covalent bonds solution. and do not ionize like bases do in solution. (Table E) ALCOHOLS ARE NOT BASES!)

o Arrhenius acids, Arrhenius bases, and salts o Given names or chemical formulas, identify (ionic compounds) are all electrolytes. An acids, bases, and salts as being electrolytes electrolyte is a substance which, when dissolved in water, forms a solution capable of o Determine the relative strength (strong or conducting an electric current (electricity). weak) of an electrolyte given information 2 Electrolytes can conduct electricity because on its ability to ionize in solution. . they ionize (break apart into ions) in a (Strong acids and strong bases are strong solution. electrolytes – Tables K and L list acids and bases in order from strongest to weakest. o The ability of a solution to conduct an electric If a salt is soluble, it is a strong electrolyte current depends on the concentration of the – Table F can be used to determine the ions in it (more ions, more conduction). solubility of different salts.)

o Recognize neutralization reactions when given the reaction equation o In the process of neutralization, an Arrhenius o Write neutralization reactions when given 3 acid and an Arrhenius base react to form a salt the reactants. (Remember that this is a . and water. double replacement reaction. Just switch Acid + Base  Salt + Water the positive ions, look up their charges and cross down the subscripts if needed. Then balance the equation.)

4 o Calculate the concentration or volume of a . o Titration is a laboratory process in which a solution, using titration data using the volume of solution with a known equation concentration is added to another solution of Ma x Va = Mb x Vb unknown concentration. Titrations are done (This equation is on Table T) to determine the concentration of the unknown solution.

o There are alternate acid-base theories. One o Give the alternate definitions of acids and 5 such theory states that the acid is a proton bases . donor (H + donor) and the base is a proton o Use this definition to explain why ammonia acceptor. is considered a base

o The acidity or alkalinity of an aqueous solution can be measured using the pH scale. o The pH scale measures the concentration of + + H /H3O in a solution. [H+] = 10–pH o Identify a solution as acidic, basic A pH of 1 means that the [H+] = 10–1 = 0.1M (alkaline), or neutral based upon the pH A pH of 3 means that the [H+] = 10–3 = value OR the relative concentrations of 0.001M H+/H O+ and OH–  Acids have pH values between 0 and 7 3 o Describe acidic, basic, and neutral solutions 6 (the stronger the acid, the lower the pH in terms of pH value and relative H+/H O+ . and the more 3 and OH– concentrations H+) o Differentiate between strong acids/bases and [H+]  [OH–] weak acids/bases given pH values or ion  Neutral solutions have a pH of 7 concentrations [H+] = [OH–]  Bases have pH values between 7 and 14 (the stronger the base, the higher the pH and the more OH–) [H+]  [OH–]

o Determine the new pH value of a solution o The pH scale is a logarithmic scale, which given the starting pH and the amount of + + means that a change of one pH unit changes increase or decrease in [H ]/[H3O ] (such as + + the concentration of H /H3O by a factor of tenfold, a hundredfold, or a thousandfold) ten Ex: A lake with an initial pH of 6 has been  tenfold = 10 times = 101 affected by acid rain. The acid rain has caused 7  hundredfold = 100 times = 102 a hundredfold change in the [H+] concentration .  thousandfold = 1000 times = 103 of the lake. What is the new pH of the lake? The exponents represent the CHANGE in pH Answer: pH = 4 o If a solution becomes more acidic, + + + + the pH , and the [H ]/[H3O ]  o Determine the amount that the [H ]/[H3O ] o If a solution becomes more basic, would increase or decrease given a certain + + the pH , and the [H ]/[H3O ]  change in pH

8 o The pH of a solution can be shown by using o Interpret changes in acid-base indicator . indicators. color o An indicator is a substance that changes color o Explain how different indicators can be depending on the concentration of used to distinguish between solutions with hydrogen/hydronium ions in a solution. different pH values o Identify appropriate indicators that can be used to show changes in pH values, such as during a titration, given starting and ending pH values

XII. NUCLEAR CHEMISTRY Knowledge Application o The stability of an isotope is based on the ratio of the neutrons and protons in the nucleus. o Usually when the ratio is not 1:1, the nucleus gets a little unstable and starts spitting out particles 1 so that it will have a more stable 1:1 ratio. . o Although most nuclei are stable, some are unstable and spontaneously emit radiation. We call these unstable isotopes radioactive isotopes, radioisotopes, or nuclides.

o Spontaneous decay (natural emission of radiation) by a nuclide (radioactive isotope) o Determine decay mode and write nuclear involves the release of particles and/or energy equations showing alpha decay, beta from the nucleus. decay, positron emission, and gamma o Each radioactive isotope has a specific decay radiation (*Remember to put radioactive mode (the kind of particle or energy it gives off emissions on the RIGHT side of the arrow 2 from its unstable nucleus) (Tables N and O!) – if something is released, it goes on the .  alpha decay: release of alpha particles right)  beta decay: release of beta particles  positron emission: release of positrons o Compare and contrast the 4 different types  gamma radiation : release of gamma rays of radiation in terms of mass, charge, o These emissions differ in mass, charge, ionizing power, and penetrating power. ionizing power, and penetrating power.

o Each radioactive isotope has a specific half-life o Calculate the initial amount, the fraction 3 (rate of decay). The half-life is the time it takes remaining, time elapsed, or the half-life of . for half of the radioisotope to decay/transmutate a radioactive isotope, given the other into something more stable). (Table N) variables

o Nuclear reactions are represented by equations that include symbols for elements and o Complete nuclear equations and predict radioactive emissions (with mass number in 4 missing particles in nuclear equations upper left and charge/atomic number in lower . Write nuclear equations given word left) o problems o These reactions show conservation of mass and charge o A change in the nucleus of an atom that o Distinguish between natural changes it from one element to another is called 5 transmutation (one reactant) and transmutation. This can occur naturally or can . artificial transmutation (two reactants) be done artificially by bombarding the nucleus given nuclear equations with high-energy particles.

o Compare and contrast fission and fusion o Types of nuclear reactions include fission and 6 reactions. fusion. Fission and fusion can be natural or . o Distinguish between fission and fusion artificial transmutations. reactions given nuclear equations

o Nuclear changes convert matter into energy (E = mc2) o Compare and contrast chemical reactions 7 Energy released during nuclear reactions is and nuclear reactions . o much greater than the energy released during o Describe benefits of using nuclear fission chemical reactions.

o There are risks and problems associated with radioactivity and the use of radioactive 8 isotopes, including: biological exposure, long- o Describe the risks and problems associated . term storage and disposal problems, and nuclear with using radioactive isotopes accidents which release radioactive materials into the environment.

o In addition to using nuclear fission for nuclear power, radioactive isotopes have other beneficial uses in medicine and industrial chemistry, including:  radioactive dating (ages of once-living things can be found from the ratio of C-14 to C-12 in the remains; ages of rocks can be found from the ratio of U-238 to Pb-206) 9  tracing chemical and biological processes (radioactive tracers can be injected into the body . and then x-rayed. The radioactive substance will show up on the x-ray and if there are problems, they can be detected easily)  detecting and treating of disease (Sr-90: diagnosing and treating bone cancer; I-131: diagnosing and treating thyroid disorders; Co-60: cancer treatment  radiation can be used to kill bacteria in foods (used with spices, meats, produce)