Mcat Chemistry

Front 1

Mass number and Atomic number of an element

Back 1

Front 2

quantum #s: electron address
n
l
m(l)
m(s)

Back 2

n: 1...infinity. aka energy level
l: 0...(n-1) subshell (e.g. s,p,d,f)
m(l): -l....+l orientation
m(s): +1/2 or -1/2 up or down

Front 3

Quantum Numbers

Back 3

Front 4

Electron Configuration exceptions

Back 4

those elements located at d4 and d9 want to make orbitals half-full or full.

e.g. Cr [Ar]4s1 3d5

Cr, Cu, Mo, Ag, Au

Front 5

Paramagnetic

Back 5

when there is unpaired electrons
odd # e- always paramagnetic

attracted to magnetic field

Front 6

Diamagnetic

Back 6

when electrons are paired at each orbital
very slight repulsion to magnetic field

even # e-....either para or diamagnetic.

Front 7

How many electrons per shell?

Back 7

2n^2 e-/shell

Front 8

alpha decay

Back 8

alpha = helium
only Z>83

Front 9

beta- decay

Back 9

n to p...because n/p >> 1

Front 10

beta+ (positron) decay

Back 10

p to n...N/Z << 1

Front 11

Electron capture

Back 11

p to n...N/Z << 1

Front 12

gamma decay

Back 12

Front 13

mass defect

Back 13

mass is less than the sum. Iron 56 has highest Nuc binding energy.
E/nucleon is proportional to ∆mass/nucleon

Front 14

Unstable Nuclei

Back 14

Odd # of protons and/or neutrons

Front 15

Stable Nuclei

Back 15

N/Z ~ 1 and Z < 20

magic #s for N or Z:
2, 8, 20, 50, 82, 126

Front 16

Effective Nuclear Charge (Zeff)

Zeff = group #

Back 16

Generally speaking, effective nuc. charge is charge felt by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus.
Example: Lithium, three protons and electron config. 1s2 2s1. The e- in 2s orbital is shielded from full attraction of the protons by the e- in the 1s orbital. Thus, Z* felt by the 2s electron should be one rather than three.

Front 17

Atomic Radius

Back 17

down and to the left.
Na > Na+
Cl- > Cl

Front 18

Ionization Energy

Back 18

up and right.
energy required to remove 1e-. 2nd e- is much harder to remove.
(Endothermic)
d5 and d10 Elements are more stable.

Front 19

Electron Affinity

Back 19

increase from left to right.
Energy change associated with gaining an e-
(Exothermic)

Front 20

Electronegativity

Back 20

up and right.
smaller radius = higher e- neg.

Front 21

Bonding

Back 21

Ionic
(strongest, meta/nonmetal, high m.p., high b.p.)
Metallic
(metal/metal, malleable, ductile)
Covalent
(nonmetal/nonmetal, lower m.p, b.p.)

Front 22

Covalent

Back 22

sharing pairs of e-
molecular: h2O
network solid: diamond, graphite

Front 23

Formal Charge

Back 23

Normal # V.E. - Actual # V.E.

Front 24

Intermolecular Forces

Back 24

Hydrogen Bonding (strongest of Inter's)
Dipoles
London Dispersion/Van der Waals

Front 25

Branching of alkanes

Back 25

Sandwiches

decreases b.p.
increases m.p.

Front 26

Longer Alkane (effects on b.p.)

Back 26

longer carbon chain increases SA, greater london dispersion force

therefore increases b.p.

Front 27

Phase Diagram

Back 27

Front 28

Phase Diagram for CO2

Back 28

1atm is lower than the triple point. This is why it sublimes (solid to gas)

Front 29

Phase Diagram for H2O

Back 29

This explains why ice floats (ice is less dense than liquid water).

Front 30

Calorimetry

Back 30

q = mc∆T

q = heat change
m = mass
c = specific heat
∆T = Tf - Ti

Front 31

Phase Transitions

Back 31

s-l-g (H>0, S>0)
fusion, sublimation, vaporization

g-l-s (H<0, S<0)
condensation, deposition, crystallization

Front 32

Ideal Gas

Back 32

Each gas molecule has no volume.
No intermolecular forces. All collisions are elastic.
KE ∝ T
1atm = 760 torr = 760 mmHg
STP = 0∘C, 1atm
1 mole gas = 22.4L

Front 33

Ideal Gas Law

Back 33

PV = nRT

Front 34

Dalton's Law of Partial Pressure

Back 34

P1 = X1 Ptot

Front 35

Graham's Law of Effustion

Back 35

Rate is proportional to velocity, lighter gas has faster veloctiy

M = molar mass

Front 36

Combined Gas Law

Back 36

P ∝ 1/V
P ∝ T
V ∝ T

Front 37

Solubility Rules

Back 37

1) All Group I metal, NH4+, and NO3- salts are soluble
2) Most Ag+, Pb2+, and Hg₂2+ salts are insoluble

Front 38

Freezing Pt. Depression

Back 38

m is number of ions, Ba(OH)2 has 3 ions when dissociated

Front 39

Vapor Pressure Depression (Raoult's Law)

Back 39

Front 40

Boiling pt. Elevation

Front 41

Boiling pt. Elevation

Back 41

Front 42

Osmotic Pressure

Back 42

water moves across membrane from (low solute) high water conc. to low conc.

pie = MRT or P = (n/V)RT

Front 43

Catalyst

Back 43

Speeds up reaction by lowering Ea.
Does not get used up. starts as reactant and ends up as product. (an intermediate is the reverse)

Front 44

Rxn Coordinate Diagram

Identify Transition State, Intermediates, delta H

Back 44

delta H < 0 because final is less than initial

Front 45

Rate Law

Back 45

only reactants in rate law.
coefficients do not determine order.
x, y exponents indicate order

Front 46

Common Ion Effect

Back 46

increasing concentration of ion will decrease solubility

Front 47

Le Chatelier's Principle

Back 47

Add something to a rxn, the rxn will adjust to regain equilibrium.

In Exothermic (T is product), increasing T will cause a shift to LEFT
increase in P = decrease in V
increase V then shift to the side with more moles of gas

Front 48

Strong Acids

Back 48

perchloric, chloric, hydrochloric....

Front 49

Strong Bases

Back 49

Front 50

Binary Acid Trend (anything with H)

Back 50

more e.n. the atom, stronger pull on H and makes it easier to give up H.
bigger the atom, longer and weaker the bond, therefore, easier to give up.

Front 51

OxyAcid Trend

Back 51

everything else being equal, more Oxygen, more acidic.

more e.n., more acidic

Front 52

[H+] of WEAK acids

Back 52

this equation is only for weak acids

Front 53

3 LAWS of Thermodynamics

Back 53

1. Conservation of Energy
2. ∆Suniv. > 0
3. S = 0 at 0K

Front 54

Systems: open, closed, isolated

Back 54

open: exchange of heat, work, matter
closed: exchange of heat and work
isolated: no exchange

Front 55

Neutral Cations

Back 55

Conjugates of strong bases e.g. group 1A metals
all other cations are acidic

Front 56

Neutral Anions/Bases

Back 56

Conjugates of strong acids e.g. Cl-, Br-, I-, NO3-, ClO4-, ClO3-

1 strong conjugate: HSO4-

almost all other anions are bases

Front 57

Henderson-Hasselbalch equation (or buffer equation)

Back 57

Front 58

Bond Enthalpies

Back 58

∆H = ∑Dbroken - ∑Dformed

D = bond

bond breaking is endo
bond making is exo
(think alphabetical b to endo, m to exo)

Front 59

Enthalpies of Formation (Hess's Law)

Back 59

products or reactants in elemental state have ZERO formation values. e.g. Cl2, O2 etc.

Front 60

∆G = ∆H - T∆S
(Get Higher Test Scores)

Back 60

remember trends:
postive ∆S helps to go spont.
negative ∆H helps to go spont.
T amplifies S

Front 61

Electrochemical Cells (for all cells)

Back 61

AN-ode OX-idation
RED-uction CAT-hode

e- always flow from anode to cathode (because e- are lost at anode, gained at cathode, cathode gains mass)

salt bridge: anions flow to anode, cations to cathode

Front 62

Galvanic/Voltaic Cell

Back 62

cathode +

anode -

Front 63

Electrolytic Cell

Back 63

cathode -

anode +

consumes energy: e- goes to neg. cathode (not spontaneous)

Front 64

Oxidation States

Back 64

1) Elements in their elemental form are in the zero oxidation state.
2) Group 1 metals are +1 and Group 2 metals are +2.
3) Hydrogen is +1 except when bonded to metals (when it’s –1).
4) The most electronegative elements get their typical ox state.
5) The last element not assigned balances the charge of the compound/ion.

Front 65

Spontaneous Rxn's

Back 65

∆G < 0
Ecell > 0
Q < K