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Kaplan mcat flashcards > Quicksheets Gen Chem > Flashcards

Quicksheets Gen Chem Flashcards

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MCAT flashcards
1
Q

how to calculate moles

A

moles = mass of sample/ molar mass

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2
Q

A =

A

A = mass number = number of protons

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3
Q

Z =

A

Z = atomic number = number of protons + neutrons

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4
Q

Planck’s quantum theory

A

energy emitted as electromagnetic radiation from matter exists in discrete bundles called quanta

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5
Q

Bohr’s model of the atom–how to calculate the energy of an electron

A

E = -Rh/n^2 Energy = -(Rydberg constant)/(principle quantum number^2) Rydberg constant = 2.18e-18J

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6
Q

equation to find the electromagnetic energy of photons

A

E = hc/λ Energy = (planck’s constant)(speed of light)/(wavelength)

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7
Q

Balmer series

A

n>2 to n = 2 Visible: 400-800 nm

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8
Q

Lyman series correspond to

A

n >1 to n = 1 UV: 90-125 nm

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9
Q

Heisenberg uncertainty principle

A

it is impossible to determine with perfect accuracy the momentum and the position of an electron simultaneously

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10
Q

principle quantum number (n)

A

the larger the integer value of n, the higher the energy level and radius of the electron’s orbit; max # of electrons in energy level n is 2n^2

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11
Q

azimuthal quantum number (l)

A

subshells l = 0,1,2,3 s,p,d,f max number of electrons that can exist w/in a subshell = 4l + 2

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12
Q

Exceptions to the octet rule (5)

A

H(2), He(2), Li(2), Be(4), B(6) and atoms found in or beyond the 3rd period

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13
Q

magnetic quantum # (ml)

A

the particular orbital w/in a sub shell where an electron is highly likely to be found at a given point in time

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14
Q

spin quantum # (ms)

A

the spin of a particle is its intrinsic angular momentum and is a characteristic of the particle, like its charge

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15
Q

Hund’s rule

A

orbitals are filled such that there are a max # of half-filled orbitals w/ parallel spins

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16
Q

polar covalent bond

A

bonding not shared equally; pulled toward more electronegative atom

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17
Q

regions of electron density: 2 Example: BeCl3 has ___ shape has ___ angle(s)

A

linear shape 180˚ angles

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18
Q

regions of electron density: 3 Example: BH3 has ___ shape has ___ angle(s)

A

trigonal planar shape 120˚ angles

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19
Q

regions of electron density: 4 Example: CH4 has ___ shape has ___ angle(s)

A

tetrahedral shape 109.5˚ angles

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20
Q

regions of electron density: 5 Example: PCl5 has ___ shape has ___ angle(s)

A

trigonal bipyramidal shape 90, 120, 180˚ angles

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21
Q

regions of electron density: 6 Example: SF6 has ___ shape has ___ angle(s)

A

octahedral shape 90˚, 180˚ angles

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22
Q

lewis acid

A

accepts electrons

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23
Q

lewis base

A

donates electrons

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24
Q

ligands

A

coordinate bonding attaches ion/molecule with metal atom

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25
Q

describe coordinate covalent bonds

A

covalent bonds in which both electrons in the bond came from the same starting atom

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26
Q

describe chelation

A

when the central cation can be bonded to the same ligand in multiple places; it generally requires large organic ligands that can double back to form a second (or even 3rd bond) with the central cation

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27
Q

chelation therapy is used to

A

often used to sequester toxic metals (lead, arsenic, mercury)

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28
Q

Describe what happens in hydrogen bonding

A

The partial positive charge of the H atom interacts with the partial negative charge located on the electronegative atoms (FON) or nearby molecules

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29
Q

describe how dipole-dipole interactions occur

A

polar molecules orient themselves such that the positive region of one molecule is close to the negative region of another molecule

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30
Q

dispersion forces

A

the bonding electrons in covalent bonds may appear to be equally shared b/t 2 atoms, but at any particular point in time they will be located randomly through the orbital permitting the unequal sharing of electrons, causing transient polarization and counterpolarization of the electron clouds of neighboring molecules, inducing the formation of more dipoles

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31
Q

a mole is…

A

the amount of substance that contains the same # of particles that are found in a 12 g sample of carbon-12

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32
Q

molecular/formula weight is measured in

A

amu per molecule (formula unit)

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33
Q

molar mass is measured in

A

grams per mole

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34
Q

What happens in combustion reactions

A

a fuel (like a hydrocarbon) is reacted with an oxidant (like oxygen) to produce an oxide and water CH4(g) + 2O2(g) -> CO2(g) + 2H2O (g)

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35
Q

What happens in combination reactions

A

two or more reactants form one product S(s) + O2(g) -> SO2(g)

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36
Q

What happens in decomposition reactions

A

a compound breaks down into 2 or more substances, usually as a result of heating or electrolysis 2HgO(s) -> 2 Hg(l) + O2(g)

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37
Q

What happens in Single-displacement rxns

A

an atom (or ion) of one compound is replaced by an atom of another element Zn(s) + CuSO4(aq) –> Cu(s) + ZnSO4(aq)

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38
Q

What happens in Double-displacement rxns

A

aka metathesis rxns elements from 2 different compounds displace each other to form two new compounds CaCl2(aq) + 2 AgNO3 (aq) –> Ca(NO3)2 (aq) + 2 AgCl(s)

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39
Q

Net ionic equations show…

A

show only the species that actually participate in the reaction (doesn’t show spectator ions) For example, in the reaction of: Zn(s) + Cu2+(aq) + SO4^2-(aq) –> Cu(s) + Zn^2+ (aq) + SO4^2-(aq) Spectator ion = SO4^2- Net ionic equation: Zn(s) + Cu^2+(aq) –> Cu(s) + Zn^2+(aq)

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40
Q

What happens in Neutralization reactions

A

These are specific double-displacement rxns that occur when an acid reacts with a base to produce a solution of a salt (and usually water) Hcl(aq) + NaOH(aq) –> NaCl(aq) + H2O(l)

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41
Q

Factors affecting reaction rates (4)

A

reactant concentrations, temperature, medium, catalysts [conc],medium, temp, catalyst

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42
Q

What do catalysts do

A

catalysts increase reaction rate w/o being consumed; do so by lowering the activation energy

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43
Q

Law of mass action states that

A

aA + bB cC + dD Kc = [C]^c[D]^d/[A]^a[B]^b Kc = equilibrium constant

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44
Q

What happens when Keq >>1

A

an equilibrium mixture contains very little of the reactants compared to the products

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45
Q

What happens when Keq << 1

A

an equilibrium mixture contains very little of the products compared to the reactants

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46
Q

What factors affect Le Chatelier’s principle?

A

stresses include concentration, pressure, volume, or temperature

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47
Q

A + B C + heat what will make it shift to the right in terms of concentration, volume, pressure, and temp?

A

Right shift = * more A or B added * C taken away * if pressure is applied or volume is reduced * If temp is reduced

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48
Q

A + B C + heat what will make it shift to the left in terms of concentration, volume, pressure, and temp?

A

Left shift = * If more C is added * If A or B is taken away * If pressure is reduced or volume is increased * If temp is increased

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49
Q

What does the law of conservation of energy say?

A

It dictates that energy can be neither created nor destroyed, but that all thermal, chemical, potential, and kinetic energies are interconvertible

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50
Q

Describe an isolated system

A

in an isolated system, there is no exchange of energy/matter with the environment

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51
Q

bomb calorimetry has to do with what kind of system

A

bomb calorimetry = a nearly isolated system

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52
Q

Describe a closed system

A

A closed system can exchange energy but not matter with the environment

53
Q

describe an open system

A

this system can exchange both energy and matter with the environment

54
Q

Describe an isothermal process

A

constant temp

55
Q

Describe an adiabatic process

A

no heat exchange

56
Q

Describe an isobaric process

A

constant pressure

57
Q

describe an isovolumetric process

A

constant volume

58
Q

isochoric process

A

constant volume

59
Q

Endothermic processes ___heat

A

endothermic = absorb heat

60
Q

Exothermic processes ____heat

A

exothermic = release

61
Q

what is constant-volume and constant pressure calorimetry used for?

A

used to indicate conditions under which the heatflow is measured

62
Q

state functions

A

Properties’ magnitudes that depend only on the initial and final states of the system, and not on the path of the change

63
Q

common state functions include (8)

A

pressure, density, temp, volume, enthalpy, internal energy, free energy, and entropy

64
Q

Enthalpy

A

(∆H) = expresses heat changes at constant pressure

65
Q

standard heat of formation

A

(∆H°f) = the enthalpy change that would occur if one mole of compound was formed directly from its elements in their standard states

66
Q

standard heat of reaction

A

(∆H°rxn) the hypothetical enthalpy change that would occur if the rxn were carried out under standard conditions

67
Q

calculation for the standard heat of reaction

A

∆H°rxn = (sum of ∆H°f of products) - (sum of ∆H°f of reactants)

68
Q

Hess’s law

A

enthalpies of reactions are additive

69
Q

bond dissociation energy

A

an average of the energy required to break a particular type of bond in one mole of gaseous molecules

70
Q

bond enthalpy

A

the standard heat of rxn that can be calculated using the values of bond dissociation energies of particular bonds

71
Q

Entropy

A

(S) it’s the measure of the distribution of energy (randomness) throughout a system

72
Q

∆Suniverse can be calculated by

A

∆Suniverse = ∆Ssystem + ∆Ssurroundings

73
Q

Gibbs free energy equation and implications

A

∆G = ∆H - T∆S -∆ G = spontaneous +∆ G = nonspontaneous ∆ G = 0 then system is in a state of equilibrium and ∆H = T∆S

74
Q

-∆H +∆S

A

spontaneous at all temps

75
Q

+∆H -∆S

A

nonspontaneous at all temps

76
Q

+∆H +∆S

A

spontaneous only at high temps

77
Q

-∆H -∆S

A

Spontaneous only at low temps

78
Q

1 atm = ? mmHg = ? torr = ? Pa

A

1 atm = 760 mmHg = 760 torr = 101,325 Pa

79
Q

STP

A

0 C/273 K, 1 atm

80
Q

standard conditions

A

25C/298K, 1 atm, 1 M concentrations

81
Q

Boyle’s law

A

PV = k P1V1 = P2V2

82
Q

Charles’s law

A

V/T = k V1/T1 = V2/T2

83
Q

Gay-Lussac’s law

A

P/T = k P1/T1 = P2/T2

84
Q

Avogadro’s Principle

A

n/V = k n1/V1 = n2/V2

85
Q

Combined gas law

A

P1V1/T1 = P2V2/T2

86
Q

ideal gas law

A

PV=nRT

87
Q

what’s the effect of decreasing the volume of a sample of gas?

A

Decreasing the volume of a sample of gas makes it behave less ideally b/c the individual gas particles are in closer proximity in a smaller volume, so they’re more likely to engage in intermolecular interactions

88
Q

Deviations due to pressure –as the pressure of a gas increases, what happens? (at moderately high pressure)

A

As the pressure of a gas increases, the particles are pushed closer and closer together. At moderately high pressure, a gas’s volume is less than would be predicted by the ideal gas law due to intermolecular attraction

89
Q

Deviations due to temp–as the temp decreases, the average velocity of the gas molecules does what? As the temp of a gas is reduced, intermolecular attraction (a) causes the gas to… At extremely low temps, what happens

A

temp of gas is reduced, gas has a smaller volume than would be predicted At extremely low temps, the volume of the gas particles themselves cause the gas to have a larger volume than predicted

90
Q

Van der Waals equation of state

A

for deviations from ideality–when a gas doesn’t closely follow the ideal gas law: (P+n^a/V^2)(V-nb) = nRT a = intermolecular attraction b = gas particles themselves

91
Q

1 mole of gas at STP = ?

A

22.4 L

92
Q

Dalton’s law of partial pressures

A

the total pressure of a gaseous mixture is equal to the sum of the partial pressures of the individual components PT = PA + PB + PC + … PA = PTXA XA = nA/nT (moles of A)/(total moles)

93
Q

kinetic molecular theory of gases says that

A

gaseous molecular behavior based on the motion of individual molecules

94
Q

average molecular speed equation

A

K = 1/2mv^2 = 3/2 KBT

95
Q

Root mean square speed equation

A

urms=sqrt:(3RT/M)

96
Q

colligative properties

A

physical properties dierived solely from the # of particles present and not the nature of those particles

97
Q

freezing point depression

A

∆Tf = iKfm =(vant Hoff factor)(Kf)(molality)

98
Q

Boiling point elevation

A

∆Tb = iKbm =(van’t Hoff factor)(Kb)(molality)

99
Q

Osmotic pressure equation

A

big pi = MRT =(Molarity)(0.0821)(Temp)

100
Q

Vapor pressure lowering is known as ____ and to calculate:

A

Vapor pressure lowering = Raoult’s law PA = XAP*A PB = XBP*B

101
Q

Occurs when gas molecules distribute through a volume by random motion

A

diffusion

102
Q

the flow of gas particles under pressure from one compartment to another through a small opening

A

effusion

103
Q

equation for diffusion/effusion

A

r1/r2 = sqrt:(m2/m1)

104
Q

Water solubility rules (4)

A
  1. all salts containing alkali metal (Group 1) or ammonium (NH4+) cations 2. all salts containing the nitrate (NO3-) or acetate (CH3COO-) anions 3. All chlorides, bromides, and iodides (except Ag+, Pb2+, and Hg2+) 4. all salts of the sulfate ion (SO42-) except Ca2+, Sr2+, Ba2+, and Pb2+
105
Q

Insolubility rules

A
  1. all metal oxides except alkali metals and CaO, SrO, BaO 6. All hydroxides except alkali metals and Ca2+, Sr2+, and Ba2+ 7. All carbonates (CO32-), phosphates (PO43-), sulfides (S2-), and sulfites (SO32-) except alkali metals and ammonium
106
Q

Calculation for the percent composition by mass

A

Mass of solute/mass of solution x 100%

107
Q

mole fraction calculation

A

of moles of compound/total # of moles in system

108
Q

molarity calculation

A

of mol of solute/L of solution

109
Q

molality calculation

A

of mol of solute/kg of solvent

110
Q

normality

A

of gram equivalent weights of solute/L of solution

111
Q

solute vs. solvent vs. solution

A

Example: Salt = solute water = solvent salt water = solution

112
Q

Arrhenius acid

A

arrhenius acid = species that produces excess H+ (protons) in an aq solution

113
Q

Arrhenius base

A

arrhenius base = species that produces excess OH- (hydroxide ions)

114
Q

Bronsted lowry acid

A

bronsted lowry acid donates protons

115
Q

Bronsted Lowry base

A

bronsted lowry base accepts protons

116
Q

lewis acid

A

lewis acid accepts electrons

117
Q

lewis base

A

lewis base donates electrons

118
Q

pH to H+ equation

A

pH = -log[H+]

119
Q

POH to OH equation

A

pOH = -log[OH-] pOH = log(1/[OH-])

120
Q

Kw equation

A

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

121
Q

pH+pOH =

A

14

122
Q

reverse reaction where the salt ions react w/ water to give back the acid or base

A

hydrolysis rxn

123
Q

oxidizing agent

A

causes another atom to undergo oxidation and it itself is reduced

124
Q

reducing agent

A

causes another atom to be reduced and it itself is oxidized

125
Q

A redox rxn in a galvanic cell has a ___∆G and is therefore ____

A

Galvanic cell = (-)∆G = spontaneous

126
Q

A redox rxn in an electrolytic cell has a ___∆G and is therefore ___

A

Electrolytic cell = (+)∆G = nonspontaneous

127
Q

EMF calculation

A

EMF = Ered,cathode- Ered,anode

128
Q

Gibbs free energy in cell calculations

A

∆G = -nFEcell

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