Ultimate GCHEM

Change in internal energy

ΔU = Q - W by system

ΔU = Q in - Q out + W on - W by

Half life

A = A0(1/2)^n

t(1/2) = 0.693/k

Energy of electron

E = -Ry/n^2

E = hc/lambda = -Ry(1/ni^2 - 1/nf^2)

Number of electrons in shell n

2n^2

Zero order reaction

Rate = k

First order reaction

Rate = k[A]

Second order reaction

Rate = k[A]^2

Rate = k[A][B]

Rate = k[B]^2

Gibbs free energy

ΔG = ΔH - TΔS

Standard free energy

ΔG0= -RT lnKeq

Keq > 1 = forward

Keq

Density of gas

d = P(molar mass)/RT = m/V

Dalton's Law of Partial Pressures

Ptotal = P1 + P2 + P3...

P1 = X1Ptotal

Boyle's Law

P1V1=P2V2

Like boil, squeezing pimple (P and V)

Charle's Law

V1/T1=V2/T2

Charlie Brown on TV (T and V)

Avogardo's Law

V1/n1 = V2/n2

V and n are directly proportional

Volume of one mole of gas at STP

22.4 L

Comparative ideal gas equation (gets rid of R)

PV/PstpVstp = nT/nstpTstp

stp = standard temperature + pressure

Molality

moles of solute/kg of solvent

Density can be helpful here

Dilution

M1V1=M2V2

Solubility Rules

See review notes

Vapor pressure depression

ΔPsoln = Xsolute*Psolvent (STP)

More solute = less surface area available for evaporation

Boiling point elevation

ΔTb = iKbm

Since boiling occurs when Pvapor = Psolvent, lowering Pvapor by adding solute increases boiling point

Freezing point depression

ΔTf = iKfm

Solute interrupts ability of solvent to form solid crystals, lowering freezing point

Osmotic pressure

π = iMRT

Pressure needed to prevent osmosis

Factors affecting solubility in water

Increase temperature: solubility of solid solute increases and solubility of gas solute decreases

Acid dissociation constant

Ka = [H+][A-]/[HA]

Base dissociation constant

Kb = [HB+][OH-]/[B]

Autoionization of water

Kw = [H+][OH-] = 1*10^-14

Log shortcut (p function)

If [anything] = m*10^-n

Then p[anything] = (n-1).(10-m)

Titration equivalence point

VaNa = VbNb

Normality equation

N = M(equivalents/mol)

Henderson-Hasselbalch equation

pH = pKa + log [A-]/[HA]

pOH = pKb + log [HB]/[B]

Electromotive force (emf)

Ecell° = Ecathode° - Eanode°

Faraday's Law of Electrolysis

It = nF

Gibbs free energy (related to emf)

ΔG = -nFEcell°

ΔG = -RTlnKeq

nFEcell° = RTlnKeq

R = 8.314

Alpha decay

Lose He: - 2 protons and neutrons

Only decay that changes atomic mass

Beta minus decay

Electron emission

One of atom's neutrons turns into a proton by emitting an electron

atomic number increases by 1

Electron capture

Beta minus decay in reverse: proton to neutron

Beta plus decay

Positron emission

atomic number decreases by 1

Gamma decay

Emission of high energy photon

no changes to atom's atomic structure

Fraction remaining after n half-lives

(1/2)^n

Fraction of nuclei decayed after n half-lives

1 - (1/2)^n

Ionic Bonds:

-Electrons leave one atom (cation) and are held by another (anion).

Covalent Bonds:

-Electrons are shared between atoms.

-Sharing can be equal (pure covalent) or unequal (polar covalent).

Metallic Bonds:

-Delocalized electrons spread over the surface of a metal.

-D Orbital.

Lewis-Dot Structures:

-Dot is a free e- in a shell.

-Line is an e- pair in a bond.

Octet Rule:

-Most atoms form bonds to achieve 8 valence e-.

-Exceptions: Hydrogen, Lithium, Beryllium, Boron.

-Expanded Octet: Phosphorus, Sulfur, Chlorine. (Group 3 or higher).

Formal Charge:

*Formal charge = valence e- minus number of pi and sigma bonds attached minus nonbonding electrons.

*# Valence e- minus (sticks + dots).

Valence Electrons:

-The electrons in the outer shell that bond or react.

Resonance Structures:

-In molecules with open p orbitals, electrons can be delocalized or shared among atoms in more than one way, creating extra stabilization.

VSEPR:

-The shape of a molecule is determined among valence e-.

Pairs:

-Bonding pairs are less repulsive due to distribution of negative charge.

-NonBonding pairs are more repulsive due to localization of negative charge.

Electronic Geometry:

-Treats bonding and nonbonding the same; does not account for angle difference.

Molecular Geometry:

-Takes into account angle differences.

Intermolecular Forces (in decreasing strength).

*Forces between different molecules.

1.) Ion-Dipole.

2.) Hydrogen Bonding.

3.) Dipole-Dipole.

4.) Dipole-Induced Dipole (Debye).

5.) London Forces / Dispersion Forces.

Principal Quantum Number: (n)

-Higher n = higher energy level = larger radius.

Angular Quantum Number: (l)

-Value can be any integer from 0 to (n-1).

*0=s, 1=p, 2=d, 3=f.

Magnetic Quantum Number: (m)

-From -L to +L.

Spin Quantum Number: (Ms)

-Can either be -1/2 or +1/2.

Hund's Rule:

-Electrons fill into a subshell such that there a maximum number of half-filled, parallel spin orbitals.

Aufbau Principle:

-Electrons fill according to the (N+L) rule such that the subshell that has the smallest sum with the smallest n fills first.

Atomic Mass: (superscript)

-(Protons + Neutrons).

Atomic Number: (subscript)

-(# of protons).

Alkali Metals:

-Group 1.

-Form +1 cations.

-Highly reactive.

Alkaline Earth Metals:

-Group 2.

-Form +2 cations.

-Somewhat reactive.

Carbon Family:

-Group 4.

-Not easily ionizable.

Nitrogen Family:

-Group 5.

-Not easily ionizable.

Oxygen Family:

-Group 6.

-Form -2 anions.

-Somewhat reactive.

Halogens:

-Group 7.

-Form -1 anions.

-Highly reactive.

Noble Gases:

-Group 8.

-Do not form ions or react.

Metals:

-Groups 1 and 2.

-Form cations.

-Ductile, malleable, shiny, conductive.

Transition Metals:

-D block.

-May be able to form cations.

-Particularly hard metals with high melting and boiling points.

Nonmetals:

-Groups 6, 7, 8.

-Brittle, poor conductors.

Metalloids/Semimetals:

-Lower left p block. (staircase).

-Mixed properties, so specific characteristics are rarely tested.

Effective Nuclear Charge: (Zeff)

-The net positive charge a valence electron feels from the nucleus.

-Increases right and up.

Atomic Radius:

-Size of the atom.

-Decreases right and up.

Electronegativity:

-Degree of attraction an atom has for electrons in a chemical bond.

-Increases right and up.

Ionization Energy:

-Energy required to completely remove an electron from an atom.

-Increases right and up.

Electron Affinity:

-Energy released when an electron is added to an atom.

-Increases right and up.

Molecules:

-Two or more atoms held together.

-Ex.) H2

Compounds:

-Require those elements to be different.

-Ex.) H2O

Molecular Mass:

-The weight of one molecule in daltons.

Molar Mass:

-The weight of one mole of molecules in grams.

Mole:

=6*10^23

Empirical Formula:

-Uses the smallest whole number ratios of atoms.

-Ex.) CH2O

Molecular Formula:

-Use the actual number of atoms.

-Ex.) C6H12O6

Percent Composition:

-Percentage of mass contributed by each element in a compound.

-Mass of (x) in formula divided by formula weight of compound.

Limiting Reactant:

-The reactant completely consumed first, halting the forward reaction.

Theoretical yield:

-Maximum yield calculated.

-We never get this through experimentation.

Actual Yield:

-The experimental yield actually measured.

Percent Yield:

-(Actual Yield divided by Theoretical Yield) * 100.

Exothermic Reaction:

-Heat given off. (exiting the system).

- (-deltaH).

Enthalpy:

-The heat of a reaction.

-The total heat content (energy) of a system.

Endothermic Reaction:

-Heat is entering system.

-(+deltaH).

Radiation:

-Transfer of heat by electromagnetic waves.

-Ex.) Sunlight, Microwave.

Conduction:

-Transfer of heat through direct contact.

Convection:

-Transfer of heat through bulk motion.

Hess's Law:

-If a reaction can be broken down into a series of steps, the enthalpy change for the overall net reaction is the sum of the enthalpies of each step.

Entropy:

-The disorder of a system.

-The energy o a system unable to do work.

Gibbs Free Energy:

-The energy of a system able to do work.

Negative Delta G:

-A spontaneous reaction.

-Will proceed forward to make products.