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UNIT 1: Atomic Structure and Properties
1.1 Moles and Molar Mass ● Avogadro’s number = 6.02 * 10^23 (1 mol of everything has this many)● Atomic mass(amu) = molar mass(g/mol) ● One unit→ another unit; use conversion factors 1.2 Mass Spectra of Elements ● Mass spectrum = graph of isotopes + abundance○ Each bar represents an isotope ● To calculate atomic mass from spectrum:∑ (isotope * relative abundance in decimal)● Mass number = protons + neutrons 1.3 Elemental Composition of Pure Substances ● To find percent composition: 1) find mass of each element 2) find mass of compound 3) (mass of each element / mass of compound) * ● Empirical formula = lowest whole number ratio of atoms vs Chemical/Molecular formula = actual number of atoms ○ Same empirical formula = same % composition ● To find empirical formula from experimental data: 1) convert all units to moles 2) divide everything by lowest number of moles 3) if needed multiply to make whole numbers ○ If given in percentages, use the percentage as the mass in grams ● For hydrate analysis problems: moles of water removed / moles of anhydrous substance 1.4 Composition of Mixtures ● Elemental analysis: finding mass of elements in a mixture to see how pure the mixture is 1.5 Atomic Structure and Electron Configuration ● Coulomb's law: F = (q1 * q2) / r^ ○ f → force/attraction; q → charges; r → distance 1.6 Photoelectron Spectroscopy ● Aufbau principle: electrons fill lower energy levels first ● Photoelectron spectroscopy: each peak represents different sublevel (ex: 1s^2, 2s^2), height of peaks represent # of electrons 1.7 Periodic Trends ● Atomic radius, ionization energy, electronegativity, electron affinity 1.8 Valence Electrons and Ionic Compounds
UNIT 2: Compound Structure and Properties
2.1 Types of Chemical Bonds ● Ionization energy is nucleus of one atom and electron of the same atom; chemical bond is the nucleus of one atom and electron of another
● Ionic (metal + nonmetal), covalent (nonmetals), metallic (metals, delocalized electrons) 2.2 Intramolecular Force and Potential Energy ● Energy curves show stable arrangement for atoms; most stable finds a balance between repulsive and attractive forces ○ Positive values = unstable; negative values = stable ● Lattice energy: energy required to separate ions in an ionic bond 2.3 Structure of Ionic Solids ● To conduct electricity a substance must have charged particles and be able to move freely 2.4 Structure of Metals and Alloys ● Properties of metals: conduct electricity, malleable, ductile● Alloys: combines 2+ metals ○ Substitutional: atoms with similar radius; substitute in alloy ○ Interstitial: atoms with different radius; smaller atom fill sup space between larger ones; stronger than metal 2.5 Lewis Diagrams ● Only valence electrons ● Bond energy increases as number of bonds increase● Central atoms are the least electronegative ● Elements in period 3 and below can have expanded octet● To draw lewis structure for molecules ○ Count electrons ○ Draw at least one bond between each element first○ Fill in remaining electrons ○ If octet not met, draw double, then triple bonds 2.6 Resonance and Formal Charge ● Length of resonance bonds are equal; either 1.5 or 1 ⅓ instead of double bond or single bond ● Resonance shown by double-sided arrow ● Formal charge = num of valence electrons - num of assigned electrons ○ Dominant structure = least number of nonzero charges; negative charges assigned to electronegative elements 2.7 VSEPR and Hybridization ● Valence Shell Electron Pair Repulsion ○ Bonds and lone pairs will arrange themselves as far apart as possible ● Geometry depends on how many electron domains (bonds / lone pairs) around central atom ○ Steric number ○ Lone pairs are more repulsive ● VSEPR CHART AND BOND ANGLES● Bond polarity
○ For a gas to behave ideally, high temp (moving faster so no time for attractions) and low pressure (so volume is insignificant) ○ Non-ideal behavior: low temp, high pressure, significant IMFs, big molecular size 3.7 Solutions and Mixtures ● Molarity = moles of solute / liters of solution 3.8 Representations of Solutions 3.9 Separation of Solutions and Mixtures ● Chromatography: uses different attractive forces in a solution ○ The more polar component has more interaction with the polar paper, and it will travel slowly ● Distillation: uses different strengths of IMFs and what effect they have on vapor pressure 3.10 Solubility ● Substances dissolve in substances with similar IMFs 3.11 Spectroscopy and the Electromagnetic Spectrum ● Microwave radiation = rotating molecules● Infrared radiation = vibrating ● UV = electrons transition to new energy levels 3.12 Properties of Photons ● Speed of light (3 * 10^8 m/s) = wavelength (m) * frequency (s^-1 or Hz) ● Energy of a photon (J) = Planck's constant (6.626 * 10^-34 J/s) * frequency (s^-1 or Hz) 3.13 Beer-Lambert Law ● Absorbance = molar absorptivity (usually a constant) * path length (usually a constant) * concentration ○ Can be simplified to absorbance = concentration● Dilution equation: M1 * V1 = M2 * V
UNIT 4: Chemical Reactions
4.1 Introduction for Reactions ● Physical: composition does not change (shape + solubility can change)● Chemical: new substance (chemical properties change) ○ Gas formation, precipitate, color change, heat/light produced 4.2 Net Ionic Equations ● Balanced molecular: show all atoms participating ● Complete ionic: show all ions in aqueous solutions● Net ionic: do not include spectator ● Solubility rules: ○ NAG SAG (nitrate, ammonium, group 1, sulfates (except PMS (lead, mercury, silver) and Ca, Sr, Ba, acetates, group 7 (expect PMS)) 4.3 Representations of Reactions 4.4 Physical and Chemical Changes
4.5 Stoichiometry ● Theoretical yield = stoichiometric calculations vs actual yield = experimental data● Percent yield = actual yield / theoretical yield 4.6 Introduction to Titration ● Equivalence point = titrant added from buret completely reacted with analyte ● Titration equation: Molarity of acid * volume of acid = molarity of base * volume of base 4.7 Types of Chemical Reactions ● Acid-base: either water and salt are products, or two types of acid-base pairs are present● Redox: oxidation numbers change ● Precipitation: mixing ions in aqueous solutions to form a solid 4.8 Introduction to Acid-Base Reactions ● Bronstead-lowry acid = proton donor; base = proton acceptor 4.9 Oxidation-Reduction (Redox) Reactions ● Half rxns only include metal oxidized or reduced
UNIT 5: Kinetics
5.1 Reaction Rates ● Rate is influenced by anything that affects the number of collisions○ Temp., volume, surface area, concentrations, catalysts ● Units for rxn rate = concentration/time 5.2 Introduction to Rate Law ● Rate law = concentration of each reactant raised to a power (order of reactant)● Rate law units based on rxn order 5.3 Concentration Changes Over Time ● Order of rxn from graph: ○ 0 order = concentration vs time is linear ○ 1st order = ln concentration vs time is linear ○ 2nd order = reciprocal concentration vs time is linear○ k = slope ● Half-life = time it takes for sample to decrease by half● Constant half-life means first order rxn 5.4 Elementary Reactions ● Elementary rxn is process in chemical rxn (single step; part of whole chemical rxn)● Substance formed then consumed = intermediates; not included in rxn ● rxns with 3+ molecules as reactants are very unlikely 5.5 Collision Model ● Molecules need proper orientation and sufficient energy to overcome activation energy 5.6 Reaction Energy Profile ● Activated complex = bonds partially formed/broken 5.7 Introduction to Reaction Mechanisms
● Amount of heat released/absorbed corresponds to coefficients in chemical equation (stoichiometry) ○ - heat treated like a product○
- heat treated like a reactant 6.7 Bond Enthalpies ● Change in energy = sum of bonds broken (reactant) + sum of bonds formed (products) 6.8 Enthalpy of Formation ● Enthalpy of formation is enthalpy change when 1 mol of compound forms from elements in their standard states ● For pure elements in standard state EOF is 0; not in standard state, EOF is not 0 ● Change in energy = sum of enthalpies of formation for products - sum of enthalpies of formation for reactants 6.9 Hess’s Law ● Reversed = change sign; total enthalpy = sum of all elementary steps
UNIT 7: Equilibrium
7.1 Introduction to Equilibrium ● Rates of forward and reverse reaction are the same ● Concentrations of products and reactants are constant, not equal 7.2 Direction of Reversible Reactions 7.3 Reaction Quotient and Equilibrium Constant ● Reaction quotient(q): concentrations of products and reactant at any given point in time○ (Concentration of product)^coefficient / (concentration of reactant)^coefficient○ Solids are liquids not included ● Equations for equilibrium constant (k) are same as q, only difference is letter ● Q will tell us whether we need to proceed forward or reverse to reach equilibrium, k will tell us if there are more products or reactants at equilibrium 7.4 Calculating the Equilibrium Constant ● K changes at different temperatures● Substitute values to calculate k ● Use ICE table to calculate equilibrium; use x for unknown values 7.5 Magnitude of the Equilibrium Constant ● More reactant = k<1; more product = k > 7.6 Properties of the Equilibrium Constant ● Reverse reaction = inverse k (note: for H, change sign); adding rxns = multiply k (note: for H, sum); multiply rxn = exponent of k (note: for H, multiply) 7.7 Calculating Equilibrium Concentrations ● When using ICE; ignore x when it’s less than 1 * 10^- 4 ○ Only remove x from denominator ○ Don't remove x if it's being added to 0
7.8 Representations of Equilibrium 7.9 Introduction to Le Châtelier’s Principle ● Changing conditions (molarity, partial pressure, temp, volume, pressure) “stresses” a system out of equilibrium and the system will try to shift to return to equilibrium ● When diluting or concentrating (adding/removing water) molarity will more significantly change on the side with more (aq) ● Increasing pressure (for gaseous systems) causes equilibrium to shift to side with least moles of gas (opposite true) ● Catalysts do not stress a system (increase rates in both directions)● When adding temp treat heat as a product or reactant ○ Changes k value (new equilibrium established) 7.10 Reaction Quotient and Le Châtelier’s Principle ● Use Q to justify shifting right or left ○ If q>k then shift to form reactants; if q ■ Before this point HQ dominates, after this point (but before the equivalence point Q- dominates, after equivalence point OH- dominates ■ Region around the half-equivalence point is called the buffer zone ● Because there is both the weak acid and conjugate base (this makes it a buffer solution ● To sketch a titration curve: ○ As strength of an acid decreases■ Starting pH increases ■ Bigger jump to buffer zone ■ Higher pH at equivalence point ● The weaker the acid, the stronger the conjugate base ● Weak base + Strong acid: ○ For example weak base (Y)○ Starting point high pH ○ Equivalence point less than 7 ■ This is because at equivalence point all Y has been neutralized, so only HY+ (conjugate acid) is present ■ Conjugate acid of a weak base is strong and will react with water according to the equation: (HY+) + H2O→ Y + (H3O+), which forms hydrogen ions, leading to an acidic solution ● NOT because the titrant acid is stronger than the base○ At half-equivalence point ■ Weak base = conjugate acid■ pOH = pKb ■ pH = 14 - pKb ■ Before this point weak base dominates, after this point conjugate acid dominates, after equivalence point H+ dominates ■ Region is also a buffer zone● To sketch a titration curve: ○ As strength of a base decreases○ Starting pH decrease ○ Bigger drop to buffer zone ○ lower pH at equivalence point ■ The weaker the base, the stronger the conjugate acid 8.6 Molecular Structure of Acids and Bases ● Stronger bonds = weaker acid (not able to dissociate) ● More electronegative atoms = stronger acid (electronegative atoms pull e- away from the H, making it less stable) 8.7 pH and pKa ● pH < pKa; acid has higher concetration
● pH > pKa; base has greater concentration 8.8 Properties of Buffers ● Mixture of acid + conjugate base ● Added acid reacts w conjugate base, added base reacts w acid; pH does not change 8.9 Henderson-Hasselbalch Equation ● pH = pKa + log (conjugate base / acid) ○ When conjugate base and acid are equal the log drops and pH = pKa● If conjugate base is greater, log is positive and pH > pKa, and opposite ○ Lesson 8. 8.10 Buffer Capacity ● Buffer capacity depends on number of moles on base and acid present○ Lower the concentrations, lower the capacity 8.11 pH and Solubility ● Common-ion effect, but the common ions are OH- and H+
UNIT 9: Thermodynamics and Electrochemistry
9.1 Introduction to Entropy ● Entropy (disorder) increases when matter/energy becomes more dispersed○ s < l < g 9.2 Absolute Entropy and Entropy Change ● Entropy = sum of entropy of products - sum of entropy of reactants 9.3 Gibbs Free Energy and Thermodynamic Favorability ● Gibbs free energy (G) shows if rxn is favorable● G = G products - G reactants ○ G < 0, favored ● G = H (enthalpy) - temp * S (entropy) 9.4 Thermodynamic and Kinetic Control ● Thermodynamically favored but does not occur fast; the process is under kinetic control (high activation energy, etc.) 9.5 Free Energy and Equilibrium ● G is -ve, k (equilibrium constant), is large (favors products)● G is 0, rxn at equilibrium 9.6 Free Energy of Dissolution 9.7 Coupled Reactions ● To make unfavored rxn happen: ○ External power source ○ Couple the unfavorable rxn w another favorable one ● 2 rxns that share a common intermediate can be coupled; produce overall favorable rxn 9.8 Galvanic (Voltaic) and Electrolytic Cells
● New conversion factor found on equation sheet
