MCAT MilesDown General Chemistry

A (top) Z (bottom) X (element)

A= mass number = protons + neutrons

Z= atomic number = number of protons

note: atomic weight = weighted average

Rutherford model

-1911

-electrons surround a nucleus

Bohr model

-1913

-described orbits in more detail

-farther orbits= more energy

-photon emitted when n goes lower, absorbed when n goes higher (n indicates level)

AHED mnemonic

Absorb light

Higher potential

Excited

Distant (from the nucleus)

Heisenberg Uncertainty

-it is impossible to know the momentum and position simultaneously

Hund's Rule

e- only double up in orbitals if all orbitals first have 1 electron

Pauli Exclusion Principle

paired e- must be +1/2 or -1/2

-spins have to be opposite directions

Avogadro's number

6.022 x 10^23 = 1 mol

Planck's (h) constant

6.626 x 10^-34 J*s

speed of light (c)

3.0 x 10^8 m/s

light energy equations

E= (hc)/gamma

E= hf

f= frequency

h= Planck's constant

c= speed of light

quantum number: n

name: principal

what it labels: e- energy level or shell number

possible values: 1,2,3...

(basically tells distance from nucleus)

notes: except for d- and f-orbitals, the shell number matches the row of the periodic table

quantum number: l

name: azimuthal

what it labels: 3D shape of the orbital

possible numbers: 0,1,2,...,n-1

notes:

0= s orbital

1= p orbital

2= d orbital

3= f orbital

4= g orbital

quantum number: mi

name: magnetic

what it labels: orbital sub-type

possible values: -n -> +n

quantum number: ms

name: spin

what it labels: electron spin

possible values: +1/2, -1/2

maximum electrons in terms of n

2n^2

maximum electrons in sub-shell

2(2l +1)

free radical

an atom or molecule with an unpaired electron

shape of s orbital

sphere

shape of p orbital

dumbbell

shape of d orbital

clover leaf

shape of f orbital

flower (8 lobes)

diamagnetic

all electrons are paired

-repelled by external magnetic field

paramagnetic

1 or more unpaired electrons

-pulled into an external magnetic field

Aufbau Principle

An electron occupies the lowest-energy orbital that can receive it

atomic orbitals on the periodic table

s- alkali and alkaline earth

p- non-metals

d- metals

f- metals

alkali metals

Group 1, 1 electron in outer level, very reactive, soft, silver, shiny, low density

-Lithium, Sodium, Potassium, Rubidium, Cesium, Francium

alkaline earth metals

metallic elements in group 2 of the periodic table which are harder than the alkali metals and are also less reactive

Berylium, magnesium, calcium, strontium, barium, radium

transition metals

Groups 3-12, 1-2 electrons in the outer energy level, less reactive than alsali-earth metals, shiny, good conductor of thermal energy and electrical current, high density

post transition metals

The are soft (or brittle), have poor mechanical strength, and have melting points lower than those of the transition metals. Being close to the metal-nonmetal border, their crystalline structures tend to show covalent or directional bonding effects, having generally greater complexity or fewer nearest neighbors than other metallic elements.

metalloids

Elements that have properties of both metals and nonmetals.

-boron, silicon, germanium, arsenic, antimony, tellurium, polonium

non-metals

Low conductivity, not ductile, not malleable, brittle, dull, gas at room temp

-carbon, nitrogen, oxygen, phosphorus, sulfur, selenium

halogens

group 17/group 7A

-flourine, chlorine, bromine, iodine, astatine, tennesine

noble gases

Group 18/group 8A

helium, neon, argon, krypton, xenon, radon, oganesson

periodic table trend: Zeff (effective nuclear charge)

-pull between the nucleus and valence electrons

-unchanged going up or down the groups

-from left to right increases across periods

-means: higher effective nuclear charge causes greater attractions to the electrons, pulling the electron cloud closer to the nucleus which results in a smaller atomic radius.

periodic table trend: ionization energy

-increases going up a group

-increases from left to right across periods

-meaning: the amount of energy required to remove an electron from an isolated atom or molecule increases following the above trends

periodic table trend: electron affinity

-gain electrons

-deltaH of the reaction < 0 when gaining electrons but EA is reported as positive values

-noble gases have no affinity for electrons (it would take energy to force an electron on them)

-increases going up a group

-increases from left to right across periods

periodic table trend: electronegativity

-force the atom exerts on an electron in a bond

-of the noble gases only Kr and Xe have an EN

-increases going up a group

-increases from left to right across periods

periodic table trend: atomic size

-increases going down a group

-increases going from right to left across periods

-this trend is basically the opposite of all others

-only this direction cations < neutral < anions

electronegativity of H

exact: 2.2

about: 2.0

electronegativity of C

exact: 2.55

about: 2.5

electronegativity of N

exact: 3.04

about: 3.0

electronegativity of O

exact: 3.44

about: 3.5

electronegativity of F

exact: 3.98

about: 4.0

covalent bond

formed via the sharing of electrons between two elements of similar EN

bond order

refers to whether the covalent bond is a single, double, or triple bond

-as bond order increases bond strength and bond energy increase but bond length decreases

nonpolar bond

delta EN < 0.5

polar bonds

delta EN is between 0.5 and 1.7

coordinate covalent bonds

a single atom provides both bonding electrons

-most often found in lewis acid-base chemistry

-lewis base: donates electrons

-lewis acid: accepts electrons

intramolecular forces

from strongest to weakest:

-hydrogen bonds

-dipole dipole

-london dispersion

note: van de walls forces are a general term that includes dipole dipole and london dispersion forces

sigma bonds and pi bonds

single bond= sigma bond

double= 1 sigma bond, 1 pi bond

triple = 1 sigma bond, 2 pi bond

bond type of according to delta EN

ionic bond

formed via the transfer of one or more electrons from an element with a relatively low IE to an element with a relatively high electron affinity

delta EN > 1.7

cation

positive +

anion

negative -

crystalline lattices

large, organized arrays of ions

formal charge

# of valence electrons - ( #nonbonding electrons + #pair of bonding electrons)

sp

e- groups around central atom: 2

bonded pairs: 2, 1

lone pairs: 0, 1

electron geometry: linear

molecule shape: linear, linear

bond angle: 180 degree

sp2

e- groups around central atom: 3

bonded pairs: 3, 2, 1

lone pairs: 0, 1, 2

electron geometry: trigonal planar

molecule shape: trig planar, bent, linear

bond angle: 120 degree

sp3

e- groups around central atom: 4

bonded pairs: 4, 3, 2, 1

lone pairs: 0, 1, 2, 3

electron geometry: tetrahedral

molecule shape: trig planar, trigonal pyramidal, bent, linear

bond angle: 109.5 degree

sp3d

e- groups around central atom: 5

bonded pairs: 5, 4, 3, 2

lone pairs: 0, 1, 2, 3

electron geometry: trigonal bipyramidal

molecule shape: trigonal bipyramdial, seesaw, T-shaped, linear

bond angle: 90 & 120 degree

sp3d2

e- groups around central atom: 6

bonded pairs: 6, 5, 4

lone pairs: 0, 1, 2, 3

electron geometry: octahedral

molecule shape: octahedral, square pyramidal, square planar

bond angle: 90 degree

electron geometry

bonded and lone pairs treated the same

molecular shape

lone pairs take up less space than a bond to another atom

equivalent mass

mass of an acid that yields 1 mole of H+ or mass of a base that reacts with 1 mole of H+

GEW (gram equivalent weight)

(molar mass)/(mol H+ or e-)

-unit is grams

equivalents =

(mass of compound)/(GEW)

-unit is grams

normality

(equivalents of solute)/(liters of solution)

or M x n (number of equivalents)

-concentration of equivalents in solution

for acids

the number of equivalents (n) is the number of H+ available from a formula unit

molarity

(normality)/(mol H+ or e-)

unit is mol/L

empirical

simplest whole-number ratio of atoms

molecular

multiple of empirical formula to show exact number of atoms of each element

combination reaction

-two or more reactants forming one product

ex: 2H2 + O2 -> 2H2O

decomposition

-single reactant breaks down

-2HgO (s) -> 2Hg + O2

combustion

-involves a fuel, usually a hydrocarbon, and O2

-commonly forms CO2 and H2O

-ex: CH4 + 2O2 -> CO2 + H2O

single displacement

-an atom/ion in a compound is replaced by another atom/ion

-Cu + AgNO3 -> Ag + CuNO3

double-displacement (metathesis)

-elements from two compounds swap places

-ex: CaCl2 + 2AgNO3 -> Ca(NO3)2 + 2AgCl

neutralization

a type of double replacement reaction

-acid + base -> salt + H2O

-HCl + NaOH -> NaCl + H2O

for elements (usually metals) that can form more than one positive ion, the charge is indicated by a Roman numeral in parentheses following the name of the element

Fe2+ -> iron(II)

Fe3+ -> iron(III)

Cu+ -> copper (I)

Cu2+ -> copper (II)

older method of naming them is: -ous and -ic to the atoms with lesser and greater charge, respectively

Fe2+ -> ferrous

Fe3+ -> ferric

Cu+ -> cuprous

Cu2+ -> cupric

monoatomic anions drop the ending of the name and add -ide

H- -> hydride

F- -> flouride

O2- -> oxide

S2- -> sulfide

N3- -> nitride

P3- -> phosphide

oxyanions= polyatomic anions that contain oxygen

more oxygen = -ate

less oxygen = -ite

NO3- -> nitrate

NO2- -> nitrite

SO4 2- -> sulfate

SO3 2- -> sulfite

in extended series of oxyanions, prefixes are also used

more oxygen= hyper (per-)

less oxygen= hypo-

ClO- -> hypochlorite

ClO2- -> chlorite

ClO3- -> chlorate

ClO4- -> perchlorate

polyatomic anions that gain H+ to for anions of lower charge add the word hydrogen or dihydrogen to the front

HCO3- -> hydrogen carbonate or bicarbonate

HSO4- -> hydrogen sulfate or bisulfate

H2PO2- -> dihydrogen phosphate

-ic

-acid name

-have more oxygen than -ous

-ous

-acid name

-have fewer oxygen than -ic

chemical kinetics

m = 0

order: zeroth order

rate law: R= k

integrated rate law: [A]= [A]0-kt

half life: t1/2 = ([A]0/2k)

units of rate constant: M/s

m = 1

order: first order

rate law: R= k[A]

integrated rate law: [A]= [A]0 x e^-kt

half life: t1/2 = (ln(2)/k)

units of rate constant: 1/s

m = 2

order: second order

rate law: R= k [A]^2

integrated rate law: 1/[A] = 1/[A]0 + kt

half life: t1/2 = (1/k[A]0)

units of rate constant: 1/M*s

Reaction Order and Michaelis-Menten Curve

-at low substrate concentrations, the reaction is approximately FIRST-ORDER

-at very high substrate concentration, the reaction approximates ZERO-order since the reaction ceases to depend on substrate concentration

hydrolysis

using water to break the bonds in a molecule

gibbs free energy (delta G)

-delta G = Ea - Ea rev

-delta G

exergonic

+delta G

endergonic

Arrhenius equation

k=Ae^(-Ea/RT)

k= rate constant

A= frequency factor

Ea = activation energy

R= gas constant = 8.314 J/mol*K

T= temp in K

Trends in Arrhenius equation

higher A = higher k

higher T = higher k

(exponent gets closer to 0 and exponent becomes less negative)

definition of rate

for aA + bB -> cC + dD

rate = (- delta [A])/ (adelta(t)) = (- delta [B])/ (bdelta(t))= (- delta [C])/ (cdelta(t)) = (- delta [D])/ (ddelta(t))

rate law

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

radioactive decay

[A]t = [A]0 x e^kt