Chemistry MCAT - Johan

acidity

increases from upper left to lower right (different!)

1. down a group, size explains: larger atom means conjugate base (cation) is more stable

2. across a period, electronegativity explains: more protons means conjugate base (cation) is more stable

hybridization

sp for 2 groups, sp2 for 3 groups, sp3 for 4 groups

lone pairs determine molecular geometry: NH3 is trigonal pyramidal (sp3), H2O is bent (sp3), XeF4 is square planar (sp3d2), SF6 is octahedral (sp3d2)

octahedral is 6 bonds, 90 degrees (sp3d2)

sp3 to sp2 conversion to attain planarity and aromaticity (4n+2 electrons resonating in a ring)

amide N can convert from sp3 to sp2 to get resonance stabilized by carbonyl O

orbitals with more s character are more stable

bond length

depends on atomic radius (which increases to lower left)

depends on bond order, more electrons shared is a stronger bond and closer bond

ATP phosphates

alpha- phosphate closest to sugar

beta- middle one

gamma- phosphate on tip

covalent bonds

metal and nonmetal share electrons

electronegativity differences creates dipoles at each bond

molecular dipole found by adding up all the bond dipoles

metallic bonds

sea of electrons delocalized between metal ions

usually S and D block elements

these compounds are conductors and malleable

coordinate covalent bond

formed between atoms with lone pairs and atoms that are electron deficient

usually between transition metals and organic compounds, where compounds donate electron pairs to coordinate with the metal

coordinate number- number of compounds coordinating with metal ions

Fe2+ can form 6 coordinate covalent bonds (that completely fill up its valence shells)

Fe2+ forms coordinate covalent bonds with hemoglobin

ionic bonds

formed between cations and anions

ions dissociate in aqueous solution, become conductor

as solids, these compounds are insulators and brittle

insulator/conductor

insulator- valence electrons tightly bound to atom

conductor- delocalized electrons, metallic bonds

intermolecular forces

pulling apart atoms is always endothermic!

4 types of IMFs:

1. ion-dipole forces- ions and polar molecule

2. dipole-dipole forces- two polar molecules, align along the molecular dipoles, H-bonding is special case

3. dipole-induced dipole forces- polar, nonpolar molecules

4. london dispersion forces- Van der Waals, temporary

5. Hydrogen bonds- align along the bond dipoles, require a donor and an acceptor, only N, O, and F can do hydrogen bonds

solvation shell

A cagelike network of solvent molecules that forms around a solute in a solution

decrease in entropy

active site

acidic and basic amino acids can undergo H-bonding and ionic interactions with the substrate

ATP has negatively charged phosphates that interact well with H, R, and K

entropy

disorder, always increasing in universe

what increases entropy (S):

1. increasing number of particles

2. increasing volume

3. increasing temperature

formation of a more organized compound or state would decrease entropy

enthalpy

breaking bonds is endothermic (dH > 0)

forming bonds is exothermic (dH < 0)

heat of formation (dHf) of elements in standard state is 0

heat of reaction (dH) = Hf products - Hf reactants

multiply by number of moles

positive dH means endothermic, heat is reactant

negative dH means exothermic, heat is product

Gibbs free energy and spontaneous reactions

dG = dH - TdS

dGo = - RTlnK

dG = dGo + RTlnQ

free energy of formation (dGf) of standard state elements is 0

spontaneous process is exergonic (dG < 0)

nonspontaneous process is endergonic (dG > 0)

examples:

combustion (-dH, +dS) is spontaneous at all temperatures

freezing (-dH, -dS) is spontaneous at low temperatures

ATP -> ADP (+dH, +dS) is spontaneous at high temperatures

bond and IMF strengths

order of bond strengths:

1. covalent

2. ionic

3. metalic

4. coordinate covalent

order of IMF strengths:

1. ion-dipole

2. dipole-dipole (H-bonds)

3. dipole-induced dipole

4. london dispersion forces

phase diagram

as pressure and temperature increase, phase change from solid to liquid to gas

triple point- all three phases coexist

critical point- above point liquid and gas are no longer distinct, becomes supercritical fluid

sublimation and evaporation line up on the diagonal

melting line is positive sloping, but for water it is negative sloping (MP decreases when pressure increases, since ice has more volume than water)

density is directly proportional to external P, indirectly proportional to external T

kinetic-molecular theory of gases

an ideal gas has:

1. no IMFs

2. particles have no volume

3. temperature is average KE

4. elastic collisions with container

high temperature, low pressure creates an ideal gas since interactions are minimized

units of pressure

1 atm = 100 kPa = 760 torr = 760 mmHg

Graham's law of diffusion/effusion

rate1/rate2 = sqrt(molar mass2/molar mass1)

heavier particles diffuse slower

reaction coordinate graph

x-axis is reaction progression

y-axis is free energy

peaks are transition states

valleys are intermediates

the highest peak is the rate limiting step

difference between reactant energy and highest peak is Ea

catalysts lower peaks, does not affect equilibrium

reaction rate

factors that increase it:

1. higher concentration of reactants

2. proper orientation, catalysts help

3. minimum energy to overcome Ea

rate constant

factors that increase it:

1. increased temperature

2. decreased activation energy

3. not affected by concentration of reactants!

rate law

rate = k[A]^x[B]^y

order is sum of exponents

determine rate law:

1. coefficients of reaction become the exponents of rate law

2. solids are not included

3. rate law determined by rate limiting step

determine rate law from data:

1. determine rate constant from any trial

2. 0th order if rate is same between two trials with diff. concentrations of a reactant

3. 1st order is rate increase is same as concentration increase of reactant

4. everything else must be 2nd order

catalytic efficiency

vmax = k_cat*[total enzyme]

catalytic efficiency = k_cat/Km

equilibrium constants

K is conc. of reactants over products, K > 1 favors products

K only changes when temperature changes

Keq > 1 means dG < 0, so spontaneous

Keq < 1 means dG > 0, so non-spontaneous

Q is reaction quotient using present concentrations state:

1. Q > K favors shift to reactants

2. Q = K is in equilibrium

3. Q < K favors shift products

why? because of the equation below:

1. dG = dGo + RTlnQ

2. dGo = -RTlnK

3. dG = 0 at equilibrium

factors that affect equilibrium

Le Chatelier's principle says system in equilibrium shifts to minimize stress:

1. change in concentration/partial pressure

2. change in pressure/volume

3. change in temperature

change in concentration or partial pressure- causes shift to reduce the change, solids do not affect equilibrium!

increase in pressure- causes shift towards side with less moles of gas

decrease in volume- causes shift towards side with less moles of gas

change in temperature- treat heat as a reactant/product

solubility equilibrium

Ksp- concentrations of products since the reactant is solid

molar solubility- moles of solid dissolved, you can plug this into Ksp equation

Qsp > Ksp means that oversaturated, precipitate forms (shift towards reactants)

common ion effect- salt solubility decreases when mixed in solution with a common ion (shift towards reactants)

opposite effect is true, salt solubility increases when mixed with solution that reacts with an ion, like an acid/base (shift towards products)

association/dissociation equilibrium

Ka = [AB]/[A][B]

Kd = [A][B]/[AB]

kinetics vs. thermodynamics

kinetics- rate, intermediate, activation energy, catalyst

faster product

thermodynamics- stability, equilibrium, spontaneity, entropy, enthalpy, free energy

more stable product

for example, free energy of ion transport does not tell you about the kinetics of a channel protein

acid/base trends

acids have more electronegative atoms

atoms without H can be acids if electron deficient

increased acidity:

1. more positive charge

2. more electronegative atom (only within a period)

3. larger atom (only within a group)

bases have less electronegative atoms

atoms without lone pairs are not basic

increased basicity:

1. more negative charge

2. less electronegative atom

3. smaller atom

acidity/basicity of salts

salts will completely dissociate in water!

acidic salt contains ion that is a weak acid

cations in group I and II are not acidic (spectators!)

basic salt contains weak base

anions like Cl, Br, or I are not basic (spectators!)

determine pH of solution

1. The acidity of an element increases as one moves to the right or down the periodic table.

2. The conjugate base of a strong acid is a very weak base, forming pH neutral solutions.

3. The conjugate base of a weak acid is a kinda weak base, pH will be higher.

amphoteric

can be both acid or base

examples:

1. amino acids

2. water

3. bicarbonate

strong acids (6) and calculating pH

HCl, HBr, HI, H2SO4, HClO4 (perchloric acid), HNO3

how to calculate pH:

1. write the concentration in scientific notation

2. take the negative of the 10 exponent

3. round down (because you ignored the actual value) and say it's somewhere in between those two values

for diprotic acids, just double the concentration, you can get a negative pH

large Ka, weak conjugate base has small Kb

weak acids and calculating pH

acetic acid, H3PO4 (phosphoric acid), H2CO3 (carbonic acid), NH4, HF

Ka = [H+]^2/[HA] from ICE table

pH = -1/2log(Ka[HA])

use same estimation technique

small Ka, strong conjugate base has large Kb

pH in solution is a bit under 7

weak acids can fully dissociate just like strong acids if another reaction removes protons to shift equilibrium to the right (like bicarbonate buffer)

buffers

conjugate acid/base pair, minimize pH change

weak acid or base (acetic acid, phosphoric acid, carbonic acid are weak acids, NH3 is weak base etc.)

both acid and base must be available to push equilibrium to account for changes in conc.

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

HH equation shows that conc. of acid/base dependent on pH

if pH < pKa, then buffer is more protonated, each difference of 1 in pH is 10-fold increase in concentration

adding water dilutes your buffer, reduces buffering capacity, but does not affect equilibrium

choose a buffer within 1 of the pH you want to maintain!

henderson-hasselbalch equation

pH = pKa + log [base]/[acid]

conc. of acid/base are dependent on pH

calculate pH of buffer solution given conc.

when [base] = [acid], pH = pKa at half-equivalence point

auto-ionization of water, pH/pOH

at standard conditions:

Kw = KaKb = 1e-14

pKa = -logKa, pKb = -logKb

pKa + pKb = 14 (pH + pOH = 14)

pH = 0 is most acidic, pOH = 0 is most basic, neutral is pH = 7

titration

sigmoidal curve, indicator changes color at equivalence point

used to determine concentration of solution, where moles added equals moles in solution to reach each equivalence point

at equivalence point:

1. compound becomes deprotonated

2. moles of acid = moles of base, meaning everything is neutralized

3. center of vertical segments

at half-equivalence point:

1. pKa = pH

2. [acid] = [base], moles of acid = 1/2 moles of base, meaning half of the acid is neutralized

4. center of horizontal segments

polyprotic titration can be done with amino acids, double the moles added to reach the second equivalence point

for titrating a strong acid, the same volume of strong or weak bases is required to neutralize, the strength will only affect the pH of equivalence point