Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

MCAT > Chemistry > Flashcards

Chemistry Flashcards

Blueprint MCAT Prep

You may prefer our related Brainscape-certified flashcards:
MCAT flashcards
1
Q

Atomic Number

A

Number of protons in an atom, determines chemical identity

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Mass Number

A

Total number of protons and neutrons in the nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Isotope

A

Same number of protons, different number of neutrons (different mass number)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Bohr model of the atom

A

Electrons orbit the nucleus in spherical shells

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What happens when electrons in an atom absorb energy?

A

When energy is absorbed, electrons are promoted to higher energy levels, farther from the nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

What happens when electrons in an atom release energy?

A

When energy is released/emitted, electrons decay from higher to lower energy levels, closer to the nucleus

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

n (quantum number)

A

Principal quantum number (corresponds to the orbital radius)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

l (quantum number)

A

Azimuthal quantum number; Denose shape and subshell identity (s, p, d, or f)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Ml (quantum number)

A

Magnetic quantum number; Denotes the orientation of an orbital within a subshell

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Ms (quantum number)

A

Describes electron spin (+/- 1/2)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Pauli Exclusion Principle

A

No two electrons in the same atom can have the same set of four quantum numbers

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Aufbau Principle

A

Lower energy orbitals fill first (watch out for exceptions, Cr, Cu, and other elements in their groups)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Hund’s Rule

A

Within a subshell, each orbital will fill with one electron before they spin pair

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

What do periods and groups correspond to in the periodic table?

A

Periods = Rows
Groups = Columns (elements in the same group often share similar properties)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Atomic Radius Periodic Trend

A

Atomic radius increases moving down and to the left along the table

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Ionization Energy, Electron Affinity, and Electronegativity periodic trends

A

Increase moving up and to the right

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

What is an isoelectronic pair?

A

An element and a cation/anion with the same electron configuration

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

First Ionization Energy

A

The energy needed to completely remove an electron from an atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Molecular Weight

A

Sum of masses of individual atoms in the molecule

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Valence Electrons

A

Electrons in outermost shell; Participate in chemical bonding

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Octet Rule

A

Atoms tend to prefer having eight valence electrons, and will form bonds to achieve this

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

Exceptions to the Octet Rule

A

Incomplete Octet (stable with < 8 Valence electrions) - H (max 2) , He (max 2), Li (stable with 2), Be (stable with 4), B (stable with 6)
Expanded Octet (stable with > 8 valence electrons) - Elements from the third period and below + atoms with odd number of electrons (radicals)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

Ionic Bonds

A

large electronegativity difference, dissociate into ions; This is the strongest kind of intramolecular bonds (intra means within)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

Covalent Bond

A

Smaller electronegativity difference, do not dissociate because atoms share electrons
Nonpolar Covalent = No or virtually no electronegativity difference
Polar Covalent = Moderate electronegativity difference (dipole moment)
Coordinate Covalent = ONe atom donates both electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

Intermolecular Forces

A

Hydrogen Bonding: Relatively strong, requires F-H, N-H, or O-H
Dipole-Dipole Forces: Weaker than H bonding; seen between molecules with fixed dipoles
London Dispersion Forces: Weakest; arise from instantaneous dipoles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

Lewis Structures

A

Depict valence electrons/bonds/lone pairs of atoms and molecules

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

Formal Charge

A

FC = VA - 1/2 BE - LPE

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

Orbital Hybridization

A

Central atom attached to 2 groups: sp
3 groups: sp2
4 groups: sp3

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

VSEPR Theory

A

Uses number of bonded atoms AND LONE PAIRS to predict molecular shape

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

Bond Angles for Molecular Geometry

A

Bent: 104.5°
Trigonal Pyramidal: 107°
Tetrahedral: 109.5°
Trigonal Planar: 120°
Linear: 180°

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q

What type of intermolecular forces contribute to hydrophobic-hydrophobic interactions?

A

London-Dispersion Forces

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
32
Q

Molecular Formula

A

Gives number of each atom present in a single molecule

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
33
Q

Empirical Formula

A

Reduces molecular formula to smallest whole-number ratio

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
34
Q

Percent Composition by Mass

A

(Mass contributed to molecule by element in question) / (total mass of molecule) x 100%

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
35
Q

Chemical Reactions

A

Involve one or more reactants changing into one or more products through the breaking and forming of bonds (the atoms stay the same - only the bonding changes)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
36
Q

Synthesis Reactions

A

Two or more reactants –> One product

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
37
Q

Decomposition Reactions

A

One reactant –> Two or more products

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
38
Q

Single Displacement Reactions

A

One element/group replaces another in a compound

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
39
Q

Double Displacement Reactions

A

Two elements/groups of two different compounds exchange places

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
40
Q

Neutralization Reactions

A

Acid + Base –> Salt + Water

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
41
Q

Combustion Reactions

A

Fuel (generally a hydrocarbon) burns in oxygen; highly exothermic; Produces CO2 and H2O

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
42
Q

Oxidation-Reduction (redox) Reactions

A

Involves the transfer of electrons (Oil Rig)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
43
Q

Balancing Reactions follows two general laws, what are they?

A

Law of Conservation of Mass: Same # of each type of atom myst be present on reactant and product sides
Law of Conservation of Charge: Net charge must be the same on both sides

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
44
Q

Stoichiometry

A

Allows us to use given quantities in a chemical reaction to find unknown ones; One common set of steps (gram of reactant –> moles of reactant –> mole ratio from reaction to find moles of desired product –> Grams of desired product)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
45
Q

Limiting Reagent

A

The reagent that is fully used in a reaction (must take into account amount present AND balanced reaction)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
46
Q

Irreversible Chemical Reaction

A

Go to completion (limiting reagent is entirely consumed)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
47
Q

Reversible Chemical Reaction

A

Marked by double arrows; proceed both forward and backward simultaneously resulting in an equilibrium state

48
Q

Equilibrium in Chemical Reactions

A

A state where forward and reverse reactions are equal

49
Q

Equilibrium Constant

A

Keq … [products] / [reactants]
Does not include solids, pure liquids, or solvents

50
Q

What values of Keq favor products? Reactants?

A

Large Keq favors products (Keq > 1)
Small Keq favors reactants (Keq< 1)

51
Q

Is Keq temperature dependent?

A

Yes, Keq is temperature dependent; temperature at standard conditions is 298 K (25°C)

52
Q

What is the reaction quotient?

A

Same equation as Keq, but can include non-equilibrium concentrations

53
Q

What values of Q indicate a reaction will proceed forward? Reverse?

A

Q < Keq: reaction will proceed forward
Q = Keq: reaction is already at equilibrium
Q > Keq: Reaction will proceed in reverse

54
Q

Le Chatlier’s Principle

A

If disturbed, an equilibrium mixture will shift to relieve stress; Shifts can result from changing concentrations, temperature, or pressure/volume; No shift results from adding a catalyst or inert gas

55
Q

Is temperature and heat the same thing?

A

No; heat involves the TRANSFER of energy

56
Q

Zeroth Law of Thermodynamics

A

For thermal equilibrium, if A = B and B = C, then A = C

57
Q

First Law of Thermodynamics

A
58
Q

Enthalpy

A

Usually expressed as ∆H; Under constant pressure, ∆H = Q

59
Q

Describe enthalpy values in endothermic and exothermic processes

A

+∆H is endothermic; Heat is required
-∆H is exothermic; Heat is released

60
Q

Hess’s Law

A

Enthalpies (∆H) of each step in a reaction are additive

61
Q

Entropy

A

(S) Relates to number of possible microstates, where +∆S (increasing disorder) is favorable according to the second law of thermodynamics

62
Q

2nd Law of Thermodynamics

A

Entropy of the universe is always increasing

63
Q

Gibbs Free Energy

A

(G)
-∆G = spontaneous (exergonic) reaction
+∆G = non-spontaneous (endergonic) reaction
∆G = ∆H - T∆S

64
Q

Thermodynamic vs. Kinetic Control

A

Thermodynamic: Forms more stable product, but may have higher Ea (and thus form more slowly); favored at high temperatures under equilibrium conditions
Kinetic: Haw lower Ea (and forms more quickly) but may be less stable; favored at low temperatures

65
Q

How do intermolecular forces affect phases?

A

Strong intermolecular interactions favor solids (vs. liquids) and liquids (vs. gasses). High temperature and low pressure favor gases (vs. solids/liquids) low temperature and high pressure favor solids

66
Q

Heat of Fusion

A

Phase change; Heat necessary to convert from solid to liquid

67
Q

Heat of Vaporization

A

Phase change; Heat necessary to convert from liquid to gas

68
Q

Describe heat and temperature during a phase change

A

At phase changes, temperature remains constant while heat is being added; Between phase changes, adding heat increases the temperature (q = mc∆T)

69
Q

Describe Phase Diagrams, triple point, critical point

A

Phase Diagrams: Generally solid –> liquid –> gas in clockwise order
Triple Point: Solid, liquid, and gas are in equilibrium
Critical Point: End of liquid-gas interface (beyond critical point matter is a supercritical fluid)

70
Q

Assumptions of kinetic molecular theory for ideal gasses

A

(1) Average kinetic energy is proportional to Temperature
(2) Particles have no volume
(3) Particles exert no forces on each other. Non-ideal behavior is found at low Temperature and high pressure

71
Q

Ideal Gas Law

A

PV = nRT

72
Q

Avagadro’s Law

A

Gas law; V/n is constant (1 mole occupies 22.4 L at STP)

73
Q

Boyle’s Law

A

Gas Law; PV = Constant (P1V1 = P2V2)

74
Q

Charles’s Law

A

Gas Law; V/T = Constant (V1/T1 = V2/T2)

75
Q

Partial Pressure of a Gas

A

The pressure that a gas in a mixture would exert if it took up the same volume by itself

76
Q

Dalton’s Law

A

Gas Law; Partial pressures of components add up to form total pressure of a mixture (Pgas = Xgas*Ptotal)

77
Q

What are the units for solutions?

A

Molarity (mol/L), molality (mol/kg), ppm (mg/L), ppt (g/L)

78
Q

Colligative Properties

A

Depend solely on number of particles in solution (vapor pressure reduction, boiling point elevation, freezing point depression, osmotic pressure)

79
Q

Solubility Constant

A

Ksp; Same principles for equilibrium constant (K) and reaction quotient (Q); High solubility constant = high solubility

80
Q

Common Ion Effect

A

Presence of one ion already in solution will decrease the solubility of a compound containing that ion

81
Q

Key Solubility Rules

A

Soluble:
(1) Alkali Metals ( Li+, Na+, K+, Rb+, Cs+, Fr+) & NH+
(2) Nitrates (NO3-), Chlorates (CLO3-), and acetate (CH3COO-)
(3) Halides (Cl-, Br-, I-) except compounds containing Ag+, Pb2+, or Hg2 2+
(4) Sulfates (SO4 2-) except compounds containing Ca2+ Sr2+, Ba2+, Ag+ or Pb2+
Insoluble:
(1) Carbonates (CO3 2-), phosphates (PO4 3-), sulfide (S2-), and sulfites (SO3 2-) except compounds containing alkali metals or NH4 +
(2) Hydroxides (OH-) and metal oxides, except compounds containing alkali metals (Ca 2+, Sr 2+, or Ba 2+)

82
Q

Activation Energy

A

Refers to the energy necessary to reach the transition state

83
Q

How can the rate of reaction be altered?

A

The rate of a reaction can be increased by increasing temperature or decreasing the activation energy (Ea)

84
Q

∆G

A

∆G < 0 = Exergonic and spontaneous
∆G > 0 = Endergonic and non-spontaneous

85
Q

What properties are thermodynamic properties?

A

∆G, ∆H, and ∆S are thermodynamic properties, independent of rate

86
Q

How do catalysts work?

A

Catalysts increase reaction rate by reducing Ea

87
Q

Enzymes

A

Enzymes are biological catalysts made of proteins

88
Q

Can enzymes alter ∆G of a reaction?

A

Catalysts (and therefore enzymes) cannot turn a nonspontaneous reaction into a spontaneous one or change any thermodynamic parameters of a reaction (∆G, ∆H, or ∆S)

89
Q

Reaction Rate

A

A reaction rate is how fast reactants are consumed and how fast products are formed. It is expressed as a decrease (-) or increase (+) in concentration per unit time, each concentration divided by the stoichiometric coefficient

90
Q

What is the rate law?

A
  • The exponents x and y must be experimentally determined. They do not reflect the stoichiometric coefficients a and b
  • This rate law reflects the initial rate
  • The units of the rate constant k can be determined algebraically. Rate is in M/s unless otherwise indicated, and concentration is in M
  • The overall order of this reaction is the sum of the exponents x and y
91
Q

Arrhenius Definition of Acids and Bases

A

Acid donates H+, base donates OH-

92
Q

Bronsted-Lowry Definition of Acids and Bases

A

Acid donates H+, base accepts H+
- When a B-L acid loses H+, it becomes its conjugate base
- When a B-L base gains H+, it becomes its conjugate acid

93
Q

Lews Definition of Acids and Bases

A

Acid accepts electron pair, base donates electron pair

94
Q

Acid Nomenclature

A
  • For acids that do not contain oxygen, use prefix “hydro-“ and suffix “-ic acid”
  • For inorganic oxyacids: named as follows, depending on number of oxygen atoms (per_____ic acid, _____ic acid, ______ous acid, hypo_____ous acid
95
Q

Do the rules for chemical equilibria apply to acid-base reactions?

A

Yes

96
Q

Kw

A

Kw = autoionization constant for water
- Kw = [H3O+][OH-] = 1 x 10^-14 at 25°C
- [H3O+] = [OH-] = 1 x 10^-7 at 25°C
- Kw is temperature dependent, but [H3O+] = [OH-] in pure water

97
Q

Ka

A

Ka = Equilibrium constant for acid dissociation; High Ka corresponds to greater dissociation/stronger acid
“Strong” Acids = Acids that fully dissociate in water to produce H3O+. Ka >1.
(HI, HBr, HCl, HNO3, H2SO4, HClO4, HClO3)

98
Q

Kb

A

Kb = Equilibrium constant for base dissociation: High Kb corresponds to greater dissociation/stronger base
- “Strong” Bases = Bases that fully ionize in water to produce OH-. Kb >1.
(Hydroxides of alkali and alkali earth metals)

99
Q

Buffers

A

Solutions that resist large changes in pH
- Include weak acid + its conjugate base (or weak base + conjugate acid)
- Physiologic example: Bicarbonate buffer system

100
Q

Polyprotic Acids

A

Contain more than one H+

101
Q

Titrations

A

Concentration of unknown solution (anylate) is found using known solution (titrant).

102
Q

Equivalence Point

A

Titrations
“Endpoint” = Full neutralization
Moles H+ = moles OH-; Found on steep segment of curve
Indicator changes color (ideal pKa indicator = pH range of equivalence point)

103
Q

Half-Equivalence Point

A

Titrations
Half-Equivalence point = half of volume required for full neutralization
- Moles acid = moles conjugate base; pH = pKa; found on flat segment of curve (plateau)

104
Q

Oxidation

A

Losing Electrons

105
Q

Reduction

A

Gaining Electrons

106
Q

OIL RIG

A

Mnemonic; OIL (Oxidation is loss of electrons); RIG (reduction is gain of electrons)

107
Q

LEO the lion says GER

A

Mnemonic; LEO (lose electrons = oxidation); GER (gain electrons reduction)

108
Q

Oxidation State

A

Basis for defining oxidation-reduction (redox) reactions

109
Q

Oxidation States Rules

A

Pure Elements: Oxidation state of 0
Ions: Overall oxidation state = charge
F: -1, most other halogens usually -1 as well, unless bonded to a more electronegative ion
H: +1, unless bonded to a more electropositive element (Ex. NaH)
O: -2, except -1 in peroxides
Alkali Metals (group 1): +1, alkaline earth metals (group 2): +2
Other Atoms: Calculated as needed to reach overall oxidation state

110
Q

Reduction Potentials

A

for a given reduction half-reaction, E* (in V) measures the driving force for a reaction. By convention these are tabulated as reduction potentials. More positive = reduction is more likely

111
Q

Galvanic/Voltaic Cells

A

The reaction is spontaneous, creating a potential difference that can be used to drive current

112
Q

Where does oxidation & reduction occur?

A

Oxidation occurs at the Anode, Reduction occurs at the cathode (think vowels and consonants)

113
Q

Concentration Cells

A

Electrode made of same material. At the beginning, concentration differences drive the redox reaction as the potential goes to zero, concentration equalizes

114
Q

Electrolytic Cells

A

Current is supplied to drive nonspontaneous redox reaction

115
Q

Rechargeable Batteries

A

Combine galvanic/voltaic and electrolytic functionality; examples include lead-acid and nickel-cadmium batteries. Voltage of a battery, which drives current, is sometimes called electromotive force (emf), but this is not a force

116
Q

Redox Titrations

A

Indicators that change color between oxidized and reduced forms, or potentiometric titrations that measure changes in potential difference. Same basic principle as acid-base titrations but measure electron transfer instead of (de)protonation

117
Q

Thermodynamic Applications of Redox Chemistry

A

Spontaneous: Ecell > 0, ▲G < 0, Keq > 1
Non-Spontaneous: Ecell < 0, ▲G > 0, Keq < 1

MCAT (6 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions