ap chem exam

buffers

• resist change in pH with small amounts of H+ or OH- added

• weak acid/conjugate base or weak base/conjugate acid

• buffer capacity: ability to react with added H+ or OH-. higher concentration of components = greater capacity

average atomic mass

decimal percentages of each isotope multiplied by isotope’s atomic mass

changes to K

• flip reaction = flip K

• double reaction = square K

exothermic reactions

• -deltaH

• energy required to break bonds in reactants must be less than energy released in forming bonds in products

• heat as product

polarity

polarizability

percent ionizations

• how much an acid dissociates compared to initial concentration

• greater the Ka = greater the percent ionization

• (concentration of dissociated compound/initial concentration) x 100

thermal equilibrium

* higher temperature decreases

mass percent

• (mass/molar mass) x 100

• (found mass of sample at certain time/mass of the sample) x 100

equivalence point

• equimolar amount of conjugate acid/base is added to base/acid.

• reaction has essentially gone to completion

• concentration of conjugate acid/base is greater than that of base/acid

empirical formua

simplest whole # ratio of atoms for a compound

molecular formula

chemical formula for compound

mass spectroscopy

mass to charge ratio of compounds

electron configuration rules

• electrons fill the lowest energy level orbital first

• no two e- can have the same spin

• e- occupy separate subshells before sharing one

photoelectron spectroscopy

measures the amount of energy electrons release

periodic trend for electronegativity

increases as you go from left to right (number of protons increases)

decreases as you go down (shielding increases)

periodic trend for atomic radius

decreases as you move from left to right (number of protons increases)

increases as you go down (shielding increases)

periodic trend for ionization energy

increases as you go left to right (number of protons increases)

decreases as you go down (shielding increases)

homogenous mixture

• uniform composition

• not chemically bonded

• 2+ substances

• e.g. salt water

heterogeneous mixture

• non-uniform composition

• not chemically bonded

• 2+ substances

ionic bonds

• between metal and nonmetals

• e- transferred

covalent bonds

• between nonmetals

• e- shared

lattice energy

energy of ionic bonds

metallic bonds

sharing of free e- between metal atoms

alloys

• compounds of different metals

• interstitial: smaller metal atoms inserted between spaces

• substitutional: similar size metal atoms substituted

hybridization

atomic orbitals fuse to form new orbitalss

formal charge

• charge of element in a molecule

• number of valence electrons of that atom and subtract number of assigned electrons in Lewis structure

  • lone pairs = 2 assigned electrons

  • bonds = 1 assigned electron

resonance

• molecules bonding structure is a combination of other possible structures

• best resonance structure is that which has total formal charge of 0/close to 0

Coulomb’s Law

• shorter distances and higher charges = stronger attractions

• F = q1xq2/r^2

  • q1 = charge of nucleus

  • q2 = charge of electron

  • r = distance between charges

strongest to weakest IMFs

1. ion-ion

2. ion-dipole: ionic compounds and liquid

3. hydrogen bonding: fluorine, oxygen, nitrogen; polar

4. dipole-dipole: between two polar molecules (polar = asymmetrical)

5. dipole-induced dipole:

6. London Dispersion Forces: exist in every sample

melting point

boiling point

viscosity

resistance of fluid (liquid or gas) to change in shape, or movement of neighboring portions relative to one another

ionic solid

• ions held in fixed positions in giant 3D lattice

• not malleable or ductile

• brittle: disrupt structure → repulsion → split solid

• high b.p. and m.p.

• low vapor pressures and volatility

• only conduct electricity when molten or in solution

volatility

tendency of substance to evaporate at normal temperatures

covalent network solids

• continuous network of covalently bonded atoms that span entire surface

• very hard

• high m.p.

• do not conduct electricity

molecular solids

• made from non-metals

• weak IMFs

• low melting points

• do not conduct electricity

• hydrogen bonding, dipole-dipole, LDFs

metallic solids

• Close packed lattice of positive atoms/ions surrounded by sea of electrons

• Good conductor of electricity and heat (closely packed)

• Metallic bond: electrostatic attraction between positive and negative charge

• Malleable and ductile

kinetic molecular theory

1. far apart

2. are in constant motion

3. elastic collisions

4. no attractions/repulsions

5. average k.e. = temperature

ideal gas law

PV = nRT

solutions

* like dissolves like

Beer’s Law

A = abc represents the change in light’s energy as it passes through a material.

photons

carry energy in waves

limiting reactant

compound that runs out during reaction, which stops it

net ionic equations

* remove spectator ions to show the species that actually interact in a reaction.

combustion reactions

hydrocarbon + O2 → H2O + CO2

redox reactions

transfer of electrons

acid-base reactions

transfer of protons

precipitation reactions

formation of insoluble solids

titrations

finding equivalence point for acid-base reactions

rates of reaction

The rate at which reactants turn into products.

rate laws

• Relates to the concentration of reactants and the reaction order.

• rate = k(concentration of reactants)

integrated rate law

time affects concentration of reactant

collision theory

• collide in right orientation

• collide with enough energy

• faster this happens, the faster reaction rate.

reaction mechanisms

elementary reactions that describe steps in reaction

rate determining step

• slowest step of reaction

• limits reaction

specific heat

energy required to raise the temperature of 1g of a substance by 1°C.

enthalpy of reaction

ΔH, the amount of heat absorbed or released by a reaction.

calorimetry

• Experimental way to measure the enthalpy of reaction

• q=mCΔT

Hess’ Law

The total enthalpy of reaction is a sum of the enthalpies for each step.

enthalpy of formation

The change in enthalpy of forming 1 mole of a compound.

bond enthalpy

Σ energy of bonds broken - Σ energy of bonds formed

equilibrium conditions

• forward rate = reverse rate

• concentrations are constant

equilibrium expression and constant

• ratio of products to reactants at equilibrium

• signified by K

reaction quotient

• ratio of products to reactants at any point in reaction

• signified by Q

solubility product

• Ratios/products of soluble compounds.

• soluble: Na, K, NH4 + , and nitrate salts, SPAN

Le Chatelier’s Principle

• add concentration: shift to other side

• dilute concentration/add volume: shift toward side with more species

• increase temperature

  • endothermic: shift toward product

  • exothermic: shift toward reactant

• decrease temperature:

  • endothermic: shift toward reactant

  • exothermic: shift toward product

• decrease volume/increase pressure: shift to side with least moles of gas

• increase volume/decrease pressure: shift to side with more moles of gas

acids

• produce H+

• H+ donors

• strong: completely dissociate into ions in water

bases

• produce OH-

• H+ acceptors

• strong: completely dissociate into ions in water

acid and base dissociation constant

• less than 1, reaction favors the reactants.

• If greater, favors products.

• signified by Ka/Kb

percent dissociation

(change in concentration/initial) x 100

Henderson-Hasselbach Equation

used to find equivalence point

titration curves

pH v volume of titrant added

equivalence point

• pH = pKa

• (HA) = (A-)

entropy

• delta S

• disorder

• The amount of entropy will always increase over time.

Gibbs Free Energy

• Available energy that can be converted into work

• Spontaneous = -ΔG = Thermodynamically favorable

voltaic cells

• spontaneous reactions

• positive cell potential

standard cell potential

potential energy difference between electrodes in volts

salt bridge

• balances charge

• anions flow to anode

• cations flow to cathode

electrolytic cells

• requires outside energy force

• I = q/t

strong acids

HCl, HBr, HClO4, HI, HNO3, H2SO4

strong bases

LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

solubility rules

* soluble
* group 1 elements
* NH4+, NO3-, salts with Cl-, Br-, or I-
* except AgCl, PbBr2, Hg2Cl2

London Dispersion Forces

• present in all molecules

• nonpolar

Dipole-Dipole Forces

* polar

hydrogen bonds

• FON

• polar

heating diagram

• use q = (moles)(delta H of fusion/vaporization) for change in temperature

• use q= mcat for phase change

sig fig rules

• non-zero digits are always sig

• zeroes in between sig digits are sig

• leading zeroes are never sig, regardless of whether or not it comes after decimal point

• zeroes are only sig if they follow decimal point

phase diagrams

• triple point: where solid, liquid, and gas all exist

• critical point: liquid and gas coexist, weird stuff happens above this

ideal gas law

• PV=nRT

• increase volume = decrease pressure

• decrease volume = increase pressure

• use when big volumes, low pressure, high temperature!

vapor pressure

* collecting gas over pressure
* Ptot - vapor pressure = partial pressure
* vapor