MCAT Chemistry

Mass Unit:

Kilogram (kg)

Distance Unit:

Meter (m)

Temperature Unit:

Kelvin (K)

Amount of a Substance Unit:

Mole (mol)

Electric Current Unit:

Ampere (A)

Time Unit:

Second (s)

Luminous Intensity Unit:

Candela (cd)

Conversion Units:

Speed of Light (c):

3.000 × 10⁸ m/s

Gas Constant (R):

8.314 J/K·mol

Avogadro’s # (N_A or L):

6.022 × 10²³/mol

Planck’s Constant (h):

6.626 x10^-34 J·s

Gravitational (G):

6.673 x10^-11 N·m²/kg²

Mass and Charge of a Neutron

Mass: ~1 amu

Charge: 0

Mass and Charge of a Proton

Mass: ~1 amu

Charge: +1

Mass and Charge of an Electron

Mass: ~0 amu

Charge: -1

Changing the subatomic particles causes:

Neutrons… Protons… Electrons…

Neutrons: New mass → isotope

Protons: New element → protons define element

Electrons: New charge → ions: anions (gain electrons (-)) and cations (lose electrons (+))

Elements of a Nuclide Symbol

Left Top: mass # (neutrons + protons)

Letter(s): element symbol

Left Bottom: atomic number

Right Top: charge

A. Cobalt-60 has more neutrons than Nickel-60

Atomic Models:

Bohr Model: n + p in nuclei & e orbit around

Quantum Mechanical Model: e are in orbitals

Quantum Mechanical Orbitals (shape, # of orbitals, # of electrons total)

s: sphere ; 1 ; 2e

p: dumbbell ; 3 ; 6e

d: no shape ; 5 ; 10e

f: no shape ; 7 ; 14e

Each orbital holds 2e and there are different levels of each

Photons Equation:

E = hf

Distance between shells

The larger the distance the more energy the higher the frequency

Ground state = lowest energy

Electron movement between shells

1st → 2nd…

5th → 4th…

1st → 2nd → 3rd… = absorb / excite a photon

5th → 4th → 3rd… = release a photon (returning to ground state)

D. Electrons transitioning from the 4th to the 1st shell

Offset of Electron Configuration Sublevels

d sublevel is offset by 1: when s and p are 4, d is 3

f sublevel is offset by 2: when s and p, f is 4

Sublevels on the Period Table

Noble Gas Shortcut

Take the noble gas BEFORE and only write the electrons that come AFTER

Ex. N: 1s²2s²2p² = [He]2p²

C. 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s²

Alkaline Earth Metal will end in #s² due to location on periodic table

Electron Configuration of Ions

Add or subtract electrons

Ex. N^3- = add 3e ; Mg²⁺ = subtract 2e

Orbital Notation Rules:

Pauli Exclusion Principle

Aufbau Rule

Hund’s Rule

Pauli Exclusion Principle: one arrow up and the other down, no 2e^- w/ same exact configuration

Aufbau Rule: start at the lowest energy level

Hund’s Rule: awkward bus rule, put one up arrow in every line (within the same shell) before putting a down arrow next to an up arrow

Electron Configuration Exemptions:

Chromium & Copper

Excited Atoms

D. 1s²2s²2p⁶3s²3p⁶

Alkaline earth ions have a +2 charge, so 2 less e- (removing the s² orbital)

C. 1s²2s²2p⁶

Periodic Table Trands

Metals: left side

Non-Metals: right side, except H

Metalloids: in between

Top: Atomic Weight

Bottom: Atomic Weight

• Atomic weight is closest to the most abundant isotope

C. There is a greater abundance of ³⁵Cl

Atomic weight of Cl (35.5) is closer to 35 than 37E

Exceptions to Octet Rule

Hydrogen (H): 2e

Boron (B): 6e

Phosphorus: 8e⁺

Sulfur: 8e⁺

Formal Charge Equation:

FC = e_valence - ½·e_bond - e_free

Calculate the Formal Charge of HCN

FC_H = 1 v.e. - ½ (2e in bond) - 0 free e^- = 1 - 1 = 0

FC_C = 4 - ½ (8) - 0 = 0

FC_N = 5 - ½ (6) - 2 = 5 - 3 - 2 = 0

Formal Charge of HCN is 0

Calculate the Formal Charge of NH⁺₄

FC_H = 0

FC_N = 5 - ½ (8) - 0 = 1

Formal Charge is 1 (reflected through + charge)

Calculate the Formal Charge of PO^3-₄

FC_o = 6 - ½ (2) - 6 = -1 (the 3 oxygens with single bonds to the P → -1 × 3 = -3)

FC_o - 6 - ½ (4) - 4 = 0 (the 1 oxygen with a double bond to P)

FC_P = 5 - ½ (10) - 0 = 0

Formal Charge is -3

Ionic Bonds:

Occurs in elements that have an electronegativity difference GREATER THAN (>) 1.7 OR metal and non-metal bond despite EN difference

Electrons are transferred from one atom to another

Very high melting point

Very high boiling point

Liquid state conducts electricity

Covalent Bonds:

Elements that have an electronegativity different of LESS THAN (<) 1.7

• Triumphed by metal and non-metal bonds

Electrons are shared between atoms

• Polar, Nonpolar, Coordinate

Lower mp than ionic

Lower bp than ionic

Bond Order: Type, Length, Strength, Contains

1: single, longest, weakest, sigma

2: double, medium, medium, sigma + pi

3: triple, shortest, strongest, sigma + 2 pi

Polar Covalent Bonds

Elements that have an electronegativity difference from 0.4 to 1.7

Electrons are unequally shared between atoms

Dipoles: p = q · r ; p = dipole moment, q= charge, r = charge separation

Lower mp than ionic, higher than nonpolar

Lower bp than ionic, higher than nonpolar

Nonpolar Covalent Bonds:

Elements that have an electronegativity difference < 0.4

Electrons are shared equally between atoms

Diatomics and Alkanes

Lower mp than polar

Lower bp than polar

Coordinate Covalent Bonds

Electrons are shared between atoms

Both electrons are denoted by one of the atoms

Hybridized characteristics of both ionic and covalent bonds

Intermolecular Forces

London Dispersion Forces: induced dipoles, size dependent (the more e^-, the stronger it is, easier to polarize)

Dipole-Dipole: polar molecules, strongest in solid and liquid

Hydrogen Bonding: results from H bonded to N, O, or F

Intermolecular Forces Strength

Hydrogen > Dipole-Dipole > London Dispersion

Which has the highest boiling point?

a. CH₄

b. C₂H₆

c. C₃H₈

d. n-C₄H₁₀

d. n-C₄H₁₀

They are all London Dispersion Forces, but d is the largest molecule

Which has the highest boiling point?

a. H₂O

b. H₂S

c. NH₃

d. PH₃

a. H₂O

H bond and stronger than NH₃ H bond

Which has the highest boiling point?

a. CH₃Cl

b. H₂O

c. n-C₅H₁₂

d. NaCl

d. NaCl

Ionic compound, so highest bp of all bonds

VSPER: Linear

Bond Angle, Hybridization, Bonds, Lone Pairs, Diagram

VSPER: Trigonal Planar

Bond Angle, Hybridization, Bonds, Lone Pairs, Diagram

VSPER: Bent

Bond Angle, Hybridization, Bonds, Lone Pairs, Diagram

VSPER: Trigonal Pyramidal

Bond Angle, Hybridization, Bonds, Lone Pairs, Diagram

VSPER: Tetrahedral

Bond Angle, Hybridization, Bonds, Lone Pairs, Diagram

VSPER: T-Shaped

Bond Angle, Hybridization, Bonds, Lone Pairs, Diagram

VSPER: Seesaw

Bond Angle, Hybridization, Bonds, Lone Pairs, Diagram

VSPER: Trigonal Bipyramidal

Bond Angle, Hybridization, Bonds, Lone Pairs, Diagram

VSPER: Octahedral

Bond Angle, Hybridization, Bonds, Lone Pairs, Diagram

Empirical Formula

Simplest whole number ratio of elements in a compound

Ex. H₂O₂ → HO

Ex. C₆H₁₂ → CH₂

Molecular Formula

Actual number of atoms of each element in a compound

Equivalents

The number of moles of a particular atom in a mole of a compound

Percent Composition by Mass Equation:

%C = grams of x / grams total

Percent Composition by Mole Equation:

% C = moles of x / moles total

Percent Yield Equation:

% Y = actual yield / theoretical yield

Limiting Reagent

The reactant that is consumed first, halting further reaction

Excess Reagent

The reactant that remains after the reaction has terminated

1. Find the moles of each element

2. Divide by the smallest number

3. Get formula from moles and find formula weight

4. Divide molecular weight by empirical formula weight

5. Multiply the molecular units by the empirical formula

Reaction Classifications

Combination: S + O₂ → SO₂

Decomposition: 2 HgO → 2 Hg + O₂

Single Displacement: Zn + CuSO₄ → Cu + ZnSO₄

Double Displacement: CaCl₂ + 2 AgNO₃ → Ca(NO₃)₂ + 2 AgCl

Rate

How fast a reaction goes

Measured in M/s

Relative Rates

Factors Affecting Rate

Temperature (increases)

Activation Energy (the higher it is, the slower the rate)

Catalyst: k = Ae^-Ea/RT ; k = rate constant, A = frequency, Ea = negative = bigger activation (smaller k), R = constant, T = temp

Rate Law

Rate = k[Reactants]^0,1,2,3,etc

NaHCO₃ + HCl → H₂O + CO₂ + NaCl

B. Warming the reaction mixture

D. Rate = k[HCOOH][C₂H₆O]

1:1 ratio, so both to the power of 1

Gibbs Free Energy when positive and negative

Positive: reaction is nonspontaneous

Negative: reaction is spontaneous

K_eq

K_eq < 1 = Reactant Favored

K_eq > 1 = Product Favored

K_eq = 1 = Equilibrium

Collision Theory of Chemical Kinetics

Reaction rate increases as rate of intermolecular collisions increases

Collisions increase or decrease due to:

Reaction concentration (Le Chatelier’s, K_eq)

Temperature (physiological limits)

Medium (correct solute)

Presence of catalysts

MEMORIZE:

Reaction Orders Shortcut:

Zeroth = Zero Dependence

First = On Dependent

Second = Squared or Two

Zeroth Order Reactions

Reactants concentration will have no effect on rate

First Order Reactions:

Rate will increase / decrease linearly to reactant concentration

Double reactant = double rate

Second Order Reactions:

Rate will increase / decrease exponentially to reactant concentration

Double reactant = quadruple rate

Each reactant will have its own rate, and the rate of the overall reaction will be the…

Sum individual reactant rates

Zero Order

First Order

Step with ONE reactant

Second Order

Step with TWO reactants

Mixed Order

B. 1st Order

1:1 ratio

C. R-S-S-R → 2 R-S

Equilibrium Constant

K = Products / Reactants

K > 1 = Product Favored

K = ~1 = Equilibrium

K < 1 = Reactant Favored

Law of Mass Action

K = [products]^coefficients / [reactants]^coefficients

ONLY include Gases and Aqueous

C. [NADH][FAD]{H⁺] / [NAD⁺] [FADH₂]

products/reactants