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MCAT Quicksheets > General Chemistry > Flashcards

General Chemistry Flashcards

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MCAT flashcards
1
Q

Calculate moles

A

= mass of sample / molar mass

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2
Q

Avogadro’s number

A

6.022 x 10^23 particles

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3
Q

Isotopes

A

Same number of protons (same atomic number) but different numbers of neutrons (different mass numbers)

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4
Q

Planck’s quantum theory

A

Energy emitted as electro-magnetic radiation from matter exists in discrete bundles called quanta

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5
Q

Energy of an electron

A

E = -Rh / n^2

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6
Q

Electromagnetic energy of photons

A

E = hc / λ

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7
Q

Balmer series

A

The group of hydrogen emission lines corresponding to transitions from upper levels n > 2 to n = 2

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8
Q

Lyman series

A

The group corresponding to transitions between upper levels n > 1 to n = 1

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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

Quantum numbers

A

1: shell (n)
- Value = n
# 2: subshell (l)
- Value = 0 to n-1
# 3: orbital (ml)
- Value = between l and -l
#4: spin (ms)
- Value = + 1/2 or - 1/2

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11
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
The maximum number of electrons in energy level n is 2n^2

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12
Q

Azimuthal quantum number (l)

A

Refers to subshells. The four subshells correspond to = 0, 1, 2, 3 and are known as s, p, d, and f
The maximum number of electrons that can exist within a subshell is given by the equation 4l + 2`

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13
Q

Periodic table trends

A

To the right: increasing effective charge, ionization energy, electronegativity, and electron affinity - decreasing atomic radius
Up the table: increasing effective charge, ionization energy, electronegativity, and electron affinity - decreasing atomic radius

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14
Q

Exceptions to octet rule

A

Stable with fewer than 8: H (2), He (2), Li (2), Be (4), B (6)

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15
Q

Magnetic quantum number (ml)

A

This specifies the particular orbital within a subshell where an electron is highly likely to be found at a given point in time

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16
Q

Spin quantum number (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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17
Q

Hund’s rule

A

Within a given subshell, orbitals are filled such that there are a maximum number of half-filled orbitals with parallel spins

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18
Q

Polar covalent bond

A

Bonding electron pair is not shared equally, but pulled toward more electronegative atom

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19
Q

Polarity of molecules

A

Depends on the polarity of the constituent bonds and on the shape of the molecule
- A molecule with nonpolar bonds is always nonpolar
- A molecule with polar bonds may be polar or nonpolar depending on the orientation of the bond dipoles

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20
Q

Linear

A

Regions of electron density = 2
Angle between electron pairs = 180

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21
Q

Trigonal planar

A

Regions of electron density = 3
Angle between electron pairs = 120

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22
Q

Tetrahedral

A

Regions of electron density = 4
Angle between electron pairs = 109.5

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23
Q

Trigonal bipyrimidal

A

Regions of electron density = 5
Angle between electron pairs = 90, 120, 180

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24
Q

Octahedral

A

Regions of electron density = 6
Angle between electron pairs = 90, 180

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25
Q

Complex ion (coordination compound)

A

A lewis acid-base adduct with a cation bonded to at least one electron pair donor (including water)

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26
Q

Ligands

A

Donor molecules and use coordinate covalent bonds

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27
Q

Chelation

A

A process where the central cation of a complex ion can be bonded to the same ligand multiple times

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28
Q

Dipole-dipole interactions

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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29
Q

Dispersion forces

A

The bonding electrons in covalent bonds may appear to be equally shared between two atoms, but at any particular point in time they will be located randomly throughout the orbital
- This permits unequal sharing of electrons, causing transient polarization and counter polarization of the electron clouds of neighboring molecules, inducing the formation of more dipoles

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30
Q

Net ionic equations

A

These types of equations are written showing only the species that actually participate in the reaction

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31
Q

Neutralization reactions

A

Acid reacts with a base to produce a solution of a salt (and, usually, water)

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32
Q

Properties of the equilibrium constant

A

If Keq»1, an equilibrium mixture of reactants and products will contain very little of the reactants compared to the products
If Keq«1, an equilibrium mixture will contain very little of the products compared to the reactants
If Keq is close to one, it is in equilibrium

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33
Q

Le Châtelier’s principle

A

Will shift to RIGHT: (1) more of reactants (2) if product is removed (3) if pressure applied or volume reduced (4) if temp is reduced
Will shift to LEFT: (2) more products (2) reactants are removed (3) if pressure is reduced or volume is increased (4) if temp is increased

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34
Q

Isolated system

A

No exchange of energy/matter with the environment; bomb calorimetry creates a nearly isolated system

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35
Q

Closed system

A

Can exchange energy but not matter with the environment

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36
Q

Open system

A

Can exchange energy and matter

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37
Q

Isothermal

A

Temp of a system remains constant

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38
Q

Adiabatic

A

No heat exchange occurs

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39
Q

Isobaric

A

Pressure of a system remains constant

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40
Q

Isovolumetric (isochoric)

A

Volume remains constant

41
Q

Endothermic

A

Reactions that absorb thermal energy

42
Q

Exothermic

A

Reactions that release thermal energy

43
Q

Endergonic

A

Reactions that are nonspontaneous

44
Q

Exergonic

A

Reactions that are spontaneous

45
Q

Enthalpy (H)

A

Used to express heat changes at constant pressure

46
Q

Standard heat of formation (ΔHf)

A

The enthalpy change that would occur if one mole of a compound was formed directly from its elements in their standard states

47
Q

Hess’s law

A

States that enthalpies of reactions are additive
- The reverse of any reaction has an enthalpy of the same magnitude as that of the forward reaction, but its sign is opposite

48
Q

Bond enthalpy

A

Sum of bonds broken - sum of bonds formed

49
Q

Entropy (S)

A

The measure of the distribution of energy (“randomness”) throughout a system

50
Q

Gibbs free energy (G)

A

Combines the two factors that affect spontaneity of a reaction - changes in enthalpy, and changes in entropy
ΔG = ΔH - TΔS

51
Q

Trends in G

A
  1. If ΔG is negative - spontaneous
  2. If ΔG is positive - nonspontaneous
  3. If ΔG is zero - the system is in a state of equilibrium
    - ΔH = TΔS
52
Q

Trends in H and S

A
  1. -ΔH and + ΔS = spontaneous at all temps
  2. +ΔH and -ΔS = nonspontaneous at all temps
  3. Both + = spontaneous only at high temps
  4. Both - = spontaneous only at low temps
53
Q

STP vs standard conditions

A

STP = (0 C or 273 K, 1 atm) - generally used for gas law calculations
Standard conditions (25 C or 298 K, 1 atm, 1 M concentrations) - measuring enthalpy, entropy, Gibbs free energy, and electromotive force

54
Q

Boyle’s law

A

PV = k or P1V1 = P2V2

55
Q

Charles’s law

A

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

56
Q

Gay-Lussac’s law

A

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

57
Q

Avogadro’s principle

A

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

58
Q

Combined gas law

A

P1V1 / T1 = P2V2 / T2

59
Q

Ideal gas law

A

PV = nRT

60
Q

Dalton’s law of partial pressure

A

PT = Pa + Pb + Pc
Pa = PT(Xa)
Where Xa = na (moles of A) / nT (total moles)

61
Q

Average molecular speed

A

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

62
Q

Freezing point depression

A

ΔTf = i(Kf)m
(m) molarity
(i) van ‘t Hoff factor for ionic compounds

63
Q

Boiling point elevation

A

ΔTb = i(Kb)m

64
Q

Osmotic pressure

A

Π = MRT

65
Q

Raoult’s law (vapor pressure lowering)

A

Pa = Xa(Poa)

66
Q

Effusion

A

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

67
Q

Diffusion and effusion formula

A

r1/r2 = sq. root (m2) / (m1)

68
Q

All salts containing group 1 and ammonium cations (NH4+) are

A

Water soluble

69
Q

All salts containing nitrate (NO3-) or acetate (CH3COO-) anions

A

Water soluble

70
Q

All chlorides, bromides, and iodides

A

Water-soluble with the exception of Ag+, Pb2+, and Hg2+

71
Q

All salts of the sulfate ion (SO4(2-))

A

Water soluble with the exception of Ca2+, Sr2+, Ba2+, and Pb2+

72
Q

All metal oxides are insoluble with the exception of:

A

Alkali metals and CaO, SrO, BaO, all of which hydrolyze to form solutions of the corresponding metal hydroxides

73
Q

All hydroxides are insoluble with the exception of:

A

The alkali metals and Ca2+, Sr2+, and Ba2+

74
Q

All carbonates, phosphates, sulfides, and sulfites are insoluble with the exception of:

A

Alkali metals and ammonium

75
Q

Percent composition by mass

A

Mass of solute / mass of solution x 100

76
Q

Molarity

A

(#) of mol of solute / liter of solution

77
Q

Mole fraction

A

(#) of mol of compound / total # of moles in system

78
Q

Molality

A

(#) of mol of solute / kg of solvent

79
Q

Normality

A

(#) of gram equivalent weights of solute / liter of solution

80
Q

Arrhenius

A

An acid is a species that produces excess H+ (protons) in an aqueous solution, and a base is a species that produces excess OH- (hydroxide ions)

81
Q

Lewis

A

An acid is an electron pair acceptor, while a base is an electron pair donor

82
Q

pH =

A

-log[H+] = log (1 / [H+])

83
Q

pOH =

A

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

84
Q

Kw =

A

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

85
Q

pH + pOH =

A

14

86
Q

Strong acids

A

HCl, HI, HBr, H2SO4, HClO3, HClO4, and HNO3

87
Q

Strong bases

A

LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, and Ba(OH)2

88
Q

Henderson-Hasselbalch equation

A

pH = pKa + log [conjugate base] / [weak acid]
pH = pKb + log [conjugate acid] / [weak base]

89
Q

Basic to acidic compounds

A

NaOH, NH3, HCO3-, F-, water, H2CO3, NH4+, HSO4-, HF, HCl

90
Q

Oxidizing agent

A

Causes another atom to undergo oxidation, and is itself reduced

91
Q

Reducing agent

A

Causes another atom to undergo reduction, and is itself oxidized

92
Q

Anode

A

(-)

93
Q

Cathode

A

(+)

94
Q

Galvanic cell

A

A redox reaction occurring here has a -ΔG

95
Q

Electrolytic cell

A

A redox reaction occurring here has a +ΔG

96
Q

Reduction potential

A

Each species is defined as the tendency of a species to acquire electrons and be reduced

97
Q

Standard reduction potentials

A

emf = Ered, cathode - Ered, anode

98
Q

Gibbs free energy, ΔG

A

ΔG = =nFEcell

MCAT Quicksheets (6 decks)

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