Chemistry Flashcards
Blueprint MCAT Prep (117 cards)
1
Q
Atomic Number
A
Number of protons in an atom, determines chemical identity
2
Q
Mass Number
A
Total number of protons and neutrons in the nucleus
3
Q
Isotope
A
Same number of protons, different number of neutrons (different mass number)
4
Q
Bohr model of the atom
A
Electrons orbit the nucleus in spherical shells
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
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
7
Q
n (quantum number)
A
Principal quantum number (corresponds to the orbital radius)
8
Q
l (quantum number)
A
Azimuthal quantum number; Denose shape and subshell identity (s, p, d, or f)
9
Q
Ml (quantum number)
A
Magnetic quantum number; Denotes the orientation of an orbital within a subshell
10
Q
Ms (quantum number)
A
Describes electron spin (+/- 1/2)
11
Q
Pauli Exclusion Principle
A
No two electrons in the same atom can have the same set of four quantum numbers
12
Q
Aufbau Principle
A
Lower energy orbitals fill first (watch out for exceptions, Cr, Cu, and other elements in their groups)
13
Q
Hund’s Rule
A
Within a subshell, each orbital will fill with one electron before they spin pair
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)
15
Q
Atomic Radius Periodic Trend
A
Atomic radius increases moving down and to the left along the table
16
Q
Ionization Energy, Electron Affinity, and Electronegativity periodic trends
A
Increase moving up and to the right
17
Q
What is an isoelectronic pair?
A
An element and a cation/anion with the same electron configuration
18
Q
First Ionization Energy
A
The energy needed to completely remove an electron from an atom
19
Q
Molecular Weight
A
Sum of masses of individual atoms in the molecule
20
Q
Valence Electrons
A
Electrons in outermost shell; Participate in chemical bonding
21
Q
Octet Rule
A
Atoms tend to prefer having eight valence electrons, and will form bonds to achieve this
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)
23
Q
Ionic Bonds
A
large electronegativity difference, dissociate into ions; This is the strongest kind of intramolecular bonds (intra means within)
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
25
Intermolecular Forces
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
26
Lewis Structures
Depict valence electrons/bonds/lone pairs of atoms and molecules
27
Formal Charge
FC = VA - 1/2 BE - LPE
28
Orbital Hybridization
Central atom attached to 2 groups: sp
3 groups: sp2
4 groups: sp3
29
VSEPR Theory
Uses number of bonded atoms AND LONE PAIRS to predict molecular shape
30
Bond Angles for Molecular Geometry
Bent: 104.5°
Trigonal Pyramidal: 107°
Tetrahedral: 109.5°
Trigonal Planar: 120°
Linear: 180°
31
What type of intermolecular forces contribute to hydrophobic-hydrophobic interactions?
London-Dispersion Forces
32
Molecular Formula
Gives number of each atom present in a single molecule
33
Empirical Formula
Reduces molecular formula to smallest whole-number ratio
34
Percent Composition by Mass
(Mass contributed to molecule by element in question) / (total mass of molecule) x 100%
35
Chemical Reactions
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)
36
Synthesis Reactions
Two or more reactants --> One product
37
Decomposition Reactions
One reactant --> Two or more products
38
Single Displacement Reactions
One element/group replaces another in a compound
39
Double Displacement Reactions
Two elements/groups of two different compounds exchange places
40
Neutralization Reactions
Acid + Base --> Salt + Water
41
Combustion Reactions
Fuel (generally a hydrocarbon) burns in oxygen; highly exothermic; Produces CO2 and H2O
42
Oxidation-Reduction (redox) Reactions
Involves the transfer of electrons (Oil Rig)
43
Balancing Reactions follows two general laws, what are they?
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
44
Stoichiometry
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)
45
Limiting Reagent
The reagent that is fully used in a reaction (must take into account amount present AND balanced reaction)
46
Irreversible Chemical Reaction
Go to completion (limiting reagent is entirely consumed)
47
Reversible Chemical Reaction
Marked by double arrows; proceed both forward and backward simultaneously resulting in an equilibrium state
48
Equilibrium in Chemical Reactions
A state where forward and reverse reactions are equal
49
Equilibrium Constant
Keq ... [products] / [reactants]
Does not include solids, pure liquids, or solvents
50
What values of Keq favor products? Reactants?
Large Keq favors products (Keq > 1)
Small Keq favors reactants (Keq< 1)
51
Is Keq temperature dependent?
Yes, Keq is temperature dependent; temperature at standard conditions is 298 K (25°C)
52
What is the reaction quotient?
Same equation as Keq, but can include non-equilibrium concentrations
53
What values of Q indicate a reaction will proceed forward? Reverse?
Q < Keq: reaction will proceed forward
Q = Keq: reaction is already at equilibrium
Q > Keq: Reaction will proceed in reverse
54
Le Chatlier's Principle
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
Is temperature and heat the same thing?
No; heat involves the TRANSFER of energy
56
Zeroth Law of Thermodynamics
For thermal equilibrium, if A = B and B = C, then A = C
57
First Law of Thermodynamics
58
Enthalpy
Usually expressed as ∆H; Under constant pressure, ∆H = Q
59
Describe enthalpy values in endothermic and exothermic processes
+∆H is endothermic; Heat is required
-∆H is exothermic; Heat is released
60
Hess's Law
Enthalpies (∆H) of each step in a reaction are additive
61
Entropy
(S) Relates to number of possible microstates, where +∆S (increasing disorder) is favorable according to the second law of thermodynamics
62
2nd Law of Thermodynamics
Entropy of the universe is always increasing
63
Gibbs Free Energy
(G)
-∆G = spontaneous (exergonic) reaction
+∆G = non-spontaneous (endergonic) reaction
∆G = ∆H - T∆S
64
Thermodynamic vs. Kinetic Control
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
How do intermolecular forces affect phases?
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
Heat of Fusion
Phase change; Heat necessary to convert from solid to liquid
67
Heat of Vaporization
Phase change; Heat necessary to convert from liquid to gas
68
Describe heat and temperature during a phase change
At phase changes, temperature remains constant while heat is being added; Between phase changes, adding heat increases the temperature (q = mc∆T)
69
Describe Phase Diagrams, triple point, critical point
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
Assumptions of kinetic molecular theory for ideal gasses
(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
Ideal Gas Law
PV = nRT
72
Avagadro's Law
Gas law; V/n is constant (1 mole occupies 22.4 L at STP)
73
Boyle's Law
Gas Law; PV = Constant (P1V1 = P2V2)
74
Charles's Law
Gas Law; V/T = Constant (V1/T1 = V2/T2)
75
Partial Pressure of a Gas
The pressure that a gas in a mixture would exert if it took up the same volume by itself
76
Dalton's Law
Gas Law; Partial pressures of components add up to form total pressure of a mixture (Pgas = Xgas*Ptotal)
77
What are the units for solutions?
Molarity (mol/L), molality (mol/kg), ppm (mg/L), ppt (g/L)
78
Colligative Properties
Depend solely on number of particles in solution (vapor pressure reduction, boiling point elevation, freezing point depression, osmotic pressure)
79
Solubility Constant
Ksp; Same principles for equilibrium constant (K) and reaction quotient (Q); High solubility constant = high solubility
80
Common Ion Effect
Presence of one ion already in solution will decrease the solubility of a compound containing that ion
81
Key Solubility Rules
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
Activation Energy
Refers to the energy necessary to reach the transition state
83
How can the rate of reaction be altered?
The rate of a reaction can be increased by increasing temperature or decreasing the activation energy (Ea)
84
∆G
∆G < 0 = Exergonic and spontaneous
∆G > 0 = Endergonic and non-spontaneous
85
What properties are thermodynamic properties?
∆G, ∆H, and ∆S are thermodynamic properties, independent of rate
86
How do catalysts work?
Catalysts increase reaction rate by reducing Ea
87
Enzymes
Enzymes are biological catalysts made of proteins
88
Can enzymes alter ∆G of a reaction?
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
Reaction Rate
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
What is the rate law?
- 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
Arrhenius Definition of Acids and Bases
Acid donates H+, base donates OH-
92
Bronsted-Lowry Definition of Acids and Bases
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
Lews Definition of Acids and Bases
Acid accepts electron pair, base donates electron pair
94
Acid Nomenclature
- 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
Do the rules for chemical equilibria apply to acid-base reactions?
Yes
96
Kw
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
Ka
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
Kb
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
Buffers
Solutions that resist large changes in pH
- Include weak acid + its conjugate base (or weak base + conjugate acid)
- Physiologic example: Bicarbonate buffer system
100
Polyprotic Acids
Contain more than one H+
101
Titrations
Concentration of unknown solution (anylate) is found using known solution (titrant).
102
Equivalence Point
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
Half-Equivalence Point
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
Oxidation
Losing Electrons
105
Reduction
Gaining Electrons
106
OIL RIG
Mnemonic; OIL (Oxidation is loss of electrons); RIG (reduction is gain of electrons)
107
LEO the lion says GER
Mnemonic; LEO (lose electrons = oxidation); GER (gain electrons reduction)
108
Oxidation State
Basis for defining oxidation-reduction (redox) reactions
109
Oxidation States Rules
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
Reduction Potentials
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
Galvanic/Voltaic Cells
The reaction is spontaneous, creating a potential difference that can be used to drive current
112
Where does oxidation & reduction occur?
Oxidation occurs at the Anode, Reduction occurs at the cathode (think vowels and consonants)
113
Concentration Cells
Electrode made of same material. At the beginning, concentration differences drive the redox reaction as the potential goes to zero, concentration equalizes
114
Electrolytic Cells
Current is supplied to drive nonspontaneous redox reaction
115
Rechargeable Batteries
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
Redox Titrations
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
Thermodynamic Applications of Redox Chemistry
Spontaneous: E*cell > 0, ▲G < 0, Keq > 1
Non-Spontaneous: E*cell < 0, ▲G > 0, Keq < 1
