Summary of Basic Chemistry Significant Tell you about the precision of a All nonzero #s and zeros between non-zero #s are significant. 708 has 3 sig figs Figures measurement If a # is less than one, count all the #s after the first non-zero #. EX: 0.0009870 has 4 sig figs. Accuracy = correct If a # is greater than 1 and no decimal point is written, count only the non- Precision = consistent zero #s and zeros in-between non-zero #s. Ex: 408000 has 3 sig figs. If a # is greater than one and a decimal point is written , count all #s. Ex: 9487.000 has 7 sig figs Adding and Subtracting Line up the decimal point. Add or subtract the numbers. Round your answer (round the answer to the to the place of the least decimal places given (the one farthest to the left) LEAST number of decimals Example: 15.00 cm 5100 cm from the givens) + 4.352 cm + 4.3 cm 19.352 = 19.35 cm 5104.3 cm = 5100 cm Multiplying and Dividing Multiply or divide the numbers. Count the total number of significant figures (Round the answer to the in each number. Your final answer should have no more significant figures LEAST # of significant figures than the lowest number of significant figures you started with of all of your givens) Example: 6.7 moles * 1.101 g/mol = 7.3767 = 7.4 g Prefixes kilo (1 km = 1000m), deci (10dm = 1m), centi (100cm = 1m), milli (1000mm = 1m) these prefixes work for any unit (ex: 1kPa = 1000Pa, 1kL = 1000L etc…) Metric Volumes : 1L = 1 dm 3 1mL = 1cm 3 Density = mass values Temperatures : 0 °C = 273K (all gas laws need to be in Kelvin!) volume Pressures : 1atm = 760 mmHg = 101.3kPa 23 1 Mole = 22.4L = 6.022 X 10 molecules = molar mass in grams (MM) density of water = 1g/mL # moles = coefficient # in a balanced equation
Subatomic Subatomic Particle Charge Mass Location Formula Particles Proton (defines the type of atom) +1 1 Nucleus = atomic number Neutron 0 1 Nucleus = atomic mass – atomic number Electron -1 0 Orbits around the nucleus = atomic number – charge Average Weighted sum of all isotopes : Example: Atomic Mass There relative abundances of Carbon-12 and carbon-13 are 98.9% and 1.19 % Σ (% Abundance) (mass) respectively. The average atomic mass is: 12*0.989 + 13*0.0119 = 12.02 amu or (do not round the final answer, units are AMU = atomic mass units) (% * mass) + (% * mass) + …. % Error Measure of accuracy (how % Error = | actual – experimental | X 100 correct your data is) actual Molar mass Synonyms: Add up the atomic masses of each atom. Pay attention to subscripts and Formula mass parenthesis. Formula weight Example: Cu(C 2H3O2)2 = Cu + 4C + 6H + 4O molecular mass = 63.546 + (4 * 12.011) + (6 * 1.0079) + (4 * 15.999) = 181.633 g/mol molecular weight Empirical (empirical formulas are the EX: 79.8% C, 20.2% H Formula LOWEST ratio of atoms) Ste p 1for C: 79.8 g (1 mol/12.011 g) = 6.64 mol C Step 1 for H: 20.2 g (1 mol/ 1.008 g) = 20.0 mol H (no diatomics!) Step 1: Calc. moles of each atom Step 2 for C: 6.64 mol C /6.64 mol = 1 C Step 2: Divide by smallest # Step 2 for H: 20.0 mol H /6.64 mol = 3 H
moles to find a mole ratio So the empirical formula is CH 3 Molecular (molecular formulas are the The molecular mass of a compound with an empirical formula of CH 3 i s 45 g. Formula UNREDUCED ratio of atoms) What is the molecular formula? Step1: Molar mass = 45 Molar mass = multiplying Empirical mass is C + 3H = (12.011 + 3*1.008) = 15 Emprical massFM factor Step 1: 45 ÷ 15 = 3
Step 2: Step 2: 3 (CH 3) = C3H9 Multiply the empirical formula by the multiplying factor
% Mass of part X 100 EX: % composition of Na in Na 3P Composition Mass of total %Na = 3(22.98) X 100 = 69.0% 99.91 +1 -2 -2 -1 Polyatomic Ammonium (NH 4 ) Sulfate (SO 4 ) Carbonate (CO 3 ) Chlorate (ClO 3 ) -1 -1 -3 ions Hydroxide (OH ) Nitrate (NO 3 ) phosphate (PO 4 ) ate always has 1 more oxygen than ite
Mole G moles Use the compound's Molar Mass (from the periodic table) Conversions Example: 1 mole of Ca 3(PO 4)2 = 3(40) + 2(31) + 8(16) = 310grams/mol How much would 0.8moles weigh? 0.8 moles x 310 g = 248grams 1 mole
Particles moles Use Avogadro's Number: (atoms, molecules, or formula units) Example: 1 mole = 6.022 x 10 23 molecules, atoms, or formula units for ANY substance How many moles of Fe would you have if you had 2.8 x 10 23 atoms 2.8 x 10 23 atoms x 1 mole = 0.46 moles 6.022 x 10 23 atoms
Liters Moles GASES ONLY, Use 1 mole = 22.4 L for any gas at STP (STP = 0 °C and 1 atm) Example: How much space would 2.5moles of diatomic chlorine (Cl 2) occupy? 2.5 moles x 22.4 L = 56L 1 mole
Multi-step Examples: Convert 10g H2O to molecules: 1 mo1 H O 6.02 X 10 23 molec OH 23 10g H 2O 2 = 3.34 x 10 molecules H 2O 18.015 g H 2O 1 mol 2OH
Convert 2.5 L H 2O to molecules: 1 mo1 OH 6.02 X 10 23 molec OH 22 2.5 L H 2O 2 = 6.7 x 10 molecules of H 2O 22.4 L 2OH 1 mol 2OH
Remember moles and molecules are NOT the same thing! Electron Rules & Laws: Alkali metal ns 1 2 Config. Alkaline Earth ns Transition = nd ? Aufbau Principle: electrons are added one at a time 5 in order of lowest energy 1 st Halogens = np Noble gases = np 6
Pauli Exclusion Principle: orbitals have a maximum of 2 e - (must be opposite spin ↑↓) Read electron configurations from the Periodic table
Hund’s Rule: (pairing rule) electrons fill sublevels p, N = 7 e - = 1s 22s 22p 3 (note that the super scripts = e -) d and f, they must fill each sublevel with one ↑ e - before any gets a paired electron ↑↓)
Formula Ionic formulas Ionic Compounds have no charge, so add subscripts so that the total + & - Writing (also called salts ) charges are equal. Made of positive cations TRICK: Switch the charges for subscripts (leave off charge and make and negative anions sure the subscripts are not divisible by anything but 1)
4+ 2- Example: Tin (IV) oxide = Sn O , so you have Sn 2O4 Since these subscripts are divisible by 2… simplify! SnO 2
Molecular formulas Use prefixes (mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca) to (covalent) determine subscripts . Note: mono is not used for the 1st element Made of two nonmetals Examples: Carbon tetrachloride = CF 4 Tri sulfur hexoxide = S 3O6 (DO NOT REDUCE covalent!)
+1 -2 -2 -1 Polyatomic Ammonium (NH 4 ) Sulfate (SO 4 ) Carbonate (CO 3 ) Chlorate (ClO 3 ) -1 -1 -3 ions Hydroxide (OH ) Nitrate (NO 3 ) phosphate (PO 4 ) ate always has 1 more oxygen than ite
Naming Ionic Compounds Name the cation. Name the anion (change ending to –ide if an element… if a Compounds (contains a metal and/or a polyatomic ion do not change the ending –ite or –ate) polyatomic ion) (If the cation can have more than 1 charge, determine charge and include as a roman numeral)
Examples: Ca 3N2 = calcium nitr ide Ca(NO 3)2 = calcium nitr ate Fe 2O3 = Iron (III) Oxide
Molecular Compounds Use prefixes to denote subscripts (mono,di,tri,tetra,penta,hexa,hepta, (two nonmetals) octa,nona,deca)(do not start the 1 st word with mono-)
Examples: CO = carbon monoxide ; N 4O8 = tetranitrogen octoxide
Naming ( H+ with and anion) Anion ending: -ide = hydro(stem)ic acid EX: HCl = hydrochloric acid Acids -ite = (stem)ous acid EX: HNO 2 = nitrous acid -ate = (stem)ic acid EX: H 3PO 4 = phosphoric acid
Balancing Sn(ClO 3)4 → SnCl 4 + O 2 You can't change the compounds at all you can only Add coefficients to Equations Sn(ClO 3)4 → SnCl 4 + 6O2 get the same number of atoms of each element on each side of the Look for common MULTIPLES
Types of Synthesis (combination) A + B AB EX: H 2 + O 2 H 2O Equations Decomposition AB A + B EX: CaO Ca + O 2
Single Replacement A + BC B + AC EX: Li + CaSO 4 Ca + Li 2SO 4
Double Replacement AB + CD CB + AD EX: CaCl 2 + NH 4NO 3 Ca(NO 3)2 + NH 4Cl
Combustion Carbon compound (C xHy) + O 2 CO 2 + H 2O this part never changes Neutralization Acid + base salt + water (type of double replacement) H2SO 4 + NaOH Na 2SO 4 + HOH a
Molarity M = mol & M 1V1 = M 2V2 Molarity = Moles of solute concentration L (dilution) Liters of solution Don’t forget to change grams to moles before using the formula for molatrity Gas Laws The Ideal Gas Law PV=nRT (R=0.0821 L*atm/mol*K or 8.31 L* kPa/Mol*K or 62.4 L*mmHg/mol*K)
STP = Combined Gas Law P1V1/T 1 = P 2V2/T 2 Kinetic Molecular Theory 0°C and 1 atm Gases can be compressed or expanded Charles’ Law V1/T 1 = V 2/T 2 (direct proportion) V ↑T↑ Gases fill up the volume of the container All gas laws Move faster with higher temperature require Boyle’s Law P1V1 = P 2V2 (indirect proportion) P ↓V↑ Constantly moving (straight paths until they KELVIN!!! bounce off – elastic collision)
Avogadro’s Law V1/n 1 = V 2/n 2 (direct proportion) V ↑n↑ NO attractive forces Less dense than solids and liquids
Gay-Lussac’s Law P1/T 1 = P 2/T 2 (direct proportion) P ↑T↑
Dalton’s Law P1+ P 2 +P 3 + P 4 + …. = P total (add up the partial pressure of each gas in the mixture) ONLY USED when there is more than one gas in the container Stoichiometry Start with # given 1mole mole ratio definition of mole = Answer definition of mole mole ratio 1mole
remember units must cancel in diagonally in order to end up with the units needed in the end
Limiting Reactant / Reagent the reactant that you run out of Excess Reactant / Reagent the reactant that you have extra
pH & pOH pH = - log[H +] if [H +] = 10 -5 the pH is 5 Acids are pH 1-6 if the pH is 9 the [H +] = 10 -9 Bases are pH 8-14 pOH = - log[OH -] if the pH is 3 the pOH is 14 - 3 = 11 Closer to 7 is more neutral if [OH -] = 0.0001 or 10 -4 the pOH is 4 pH + pOH = 14 if the pOH is 2the [OH -] = 10 -2 or 0.01 Acids and Arrhenius: Properties of acids & bases Titrations Bases Acids have H + ; Bases have OH - BOTH are electrolytes (conduct electricity) Using a neutralization reaction The more it dissociates (breaks into ions) the between acids and bases to Bronsted-Lowry: stronger the electrolyte / acid / base determine the molarity of either Acids are proton donors; A = sour ; B = bitter the acid or the base. Bases are proton receivers A = “wet” ; B = slippery A = corrosive ; B = caustic Ma V a = Mb V b A = litmus turns red ; B = litmus turns blue C a C b coef 2 Equilibrium Equilibrium means Keq = [product] EXAMPLE : K eq = [NO] coef Constant Both and RATES are the same [reactants] [N 2][O 2] Keq (in a closed system) N 2 (g) + O 2 (g) ↔ 2NO(g) Cross off (ignore) solids and liquids; keep gases and aqueous Le Chatelier’s Equilibrium will shift to counter An increase in concentration or heat ↑ will cause a shift AWAY from what was ↑ Principle the stress applied to it A decrease in concentration or heat ↓will cause a shift TOWARDS it A decrease in volume will shift towards the side with LESS moles of gas An increase in volume will shift towards the side with MORE moles of gas (recall that volume is INVERSE to pressure; volume ↑ pressure ↓)
Heat Specific heat of H 2O = ∆T = ∆Tfinal -∆Tinitial Calories or ° 1 cal/g C Within a phase Heat (q) = (mass) x (specific heat) x ( ∆T) = mC p∆T Joules or 4.18 J = 1 cal Changing phase
Hf (melting) Heat (q) = mH Specific heat defined as f Hv (evaporation) Heat (q) = mH amount of energy needed to v
change 1 gram by 1 °°°C Oxidation & REDOX reactions Oxidized = lost e- and increase in oxidation number (gaining oxygen) Reduction L oss of Reduction = gain e- and decrease in oxidation number (loss of oxygen) Electrons is AGENTS are opposite… Oxidation The reactant that was oxidized becomes the reducing agent Says The reactant that was reduced becomes theoxidizing agent Gain of MUST look at oxidation numbers! Electrons is Reduction EX : Al + NaNO 3 Al(NO 3)3 + Na Al is oxidized (red agent) 0 +1+5-2 +3+5-2 0 Na is reduced (ox agent) ½ Reactions Ox ½ = Al 0 Al +3 + 3 e- electrons on the right of the arrow e - Red ½ = Na +1 + 1e- Na 0 electrons on the left of the arrow e - Scientists Models: Discoveries & Experiments: Idea of indivisible atom (Democritus) Thomson = cathode ray tube experiment = electron charge Billiard ball (Dalton) Millikan = oil drop experiment = mass to charge ratio of electron Plum pudding (Thomson) Rutherford = Gold foil experiment = atom’s dense nucleus Nuclear (Rutherford) Periodic Tables: Planetary (Bohr and Plank) Mendeleev = 1 st PT = based on increasing atomic mass Quantum (Schrodinger, Heisenberg – Uncertainty Mosely = modern = based on increasing atomic number Principle, deBroglie – Wave theory) Radioactivity 3 types of radiation: HALF LIFE = the amount of time needed to decay a radioactive & Half Life Alpha α = weakest = He nucleus (2 p + + 2n 0) = substance to half the mass. Shielded by paper If 4000 g of substance has a half life of 3 hours. How much Beta β = moderate = electrons = Shielded by remains after 12 hours? aluminum Gamma δ / γ = strongest = EMR rays = Shielded 12 /3 = 4 half lives ( ) take place in 12 hours by lead 4000 2000 1000 500grams
Factors Affecting An increase in any of the following Catalysts – chemical that lowers the activation energy Reaction Rates will increase the rate and speeds up the reaction Temperature, Concentration, Surface area Periodic Table Atomic Mass = proton PLUS neutrons (Basics) This number is a weighted average of the isotopes
How to read it Atomic Number = proton number In atoms protons = electrons they are neutral In ions protons ≠ electrons they are charged
Neutrons = atomic mass – atomic number the key above is given on the SOL periodic table
Isotopes Isotopes have same proton number , but different neutron number and this changes the mass
Another way for writing this isotope of Krypton is Kr – 94
Periodic Table Alkali Metals: ns 1 (Groups) (+1 oxidation; 1 valence electron) Most reactive of the metals; very soft; explodes in water to make H 2 ; oxidizes very easily Alkaline Earth Metals: ns 2 (+2 oxidation; 2 valence electrons) Moderately reactive Transition Metals: nd ? Forms alloys easily; has varying oxidation states Halogens: np 5 (-1 oxidation; 7 valence electrons) Most reactive of the nonmetals; deadly gases, diatomic Periods are ROWS ↔ Nobel Gases: np 6 (NO oxidation; 8 valence electrons except He = 2) Groups or Families are COLUMNS ↕ “inert” = not reactive due to full outer shell (octet rule) Metals vs Metals (LEFT of zig zag stair step line) Nonmetals (RIGHT of zig zag stair step line) Nonmetals Malleable (bend easily); Ductile (stretch easily) Brittle, break and shatter easily Conductive (conduct electricity easily) NonConductive (do NOT conduct electricity easily) Lustrous (shiny) Dull (not shiny) Form + charged ions (cations) by losing electrons Form – charged ions (anions) by gaining electrons Periodic Table Ionization Energy: (Trends) Energy needed to remove an electron (metals have lowest energy)
Electronegativity: Ability to attract electrons when bonding (Fluorine has the highest value)
Atomic Radii: Size of atom (metals are larger than nonmetals)
Reactivity: How well it reacts (Fr is the most reactive metal; F is the most reactive nonmetal)
Shielding Effect: The effect of filled energy levels (stonger shielding as you move down a group) Lewis Dot Lewis dot carbon Diagrams diagrams use dioxide water # valence electrons Heating curve Triple Point Relationships Curve inverse
Boyles law (pressure and volume)
direct
Charles (Volume and Temp) GayLussacs (Pressure and Temp) Vapor Pressure Curve VSEPR – Geometric shapes Valence Shell Electron Repulsion Theory
Linear X1Y2 Bent (X has no extra valence electrons) X1Y2 (X has extra valence electrons that cause the bent shape)
Trigonal Planar Tetrahedral X1Y3 X Y Nonvolatile – (X has no extra valence electrons) 1 4 high boiling point (X has no extra valence electrons) Look at Lewis Dot diagrams to help determine the shape - Volatile – low boiling point Not shown is Trigonal Pyramidal (NH 3) where N has extra ve which repell Exothermic = gives Solubility: off heat Solutes dissolve in Solvents A C + Heat “Like dissolves like”
Supersaturated ABOVE the line
Saturated ON the line Endothermic = Unsaturated A+heat C UNDER the line
Solids increase with temp
Gases decrease with temp
Graduated Mortar + Pestle Safety FIRST: Volumetric Spilled chemical – rinse off cylinder Flask (crushing) (measure) Diluting acid - Add acid TO water (Make a solution) Evaporating dish (evaporating) MSDS – Material Safety Data Sheet (read them carefully ALL the info is Funnel in them) (filtering ))