Chemistry Flashcards
1
Q
Atomic number
A
of protons, Z
2
Q
Atomic mass (mass number)
A
of protons + neutrons, A
3
Q
Isotopes
A
Structures of the same element that differ in the number of neutrons
4
Q
Atomic weight
A
Weighted average of all isotopes
5
Q
Plank’s constant
A
E = hf, h= 6.626 x 10^-34 J·s, Explain’s how matter releases energy as em radiation in quanta
6
Q
Avagadro’s number
A
6.6 x 10^23 molecules/mol
7
Q
Principal quantum number
A
n, it describes the size of the orbital (energy level). The max # of e- is 2n^2
8
Q
Azimuthal quantum number
A
l, the shape and # of subshells. Subshells 0-3 are designated spdf
9
Q
Magnetic quantum number
A
ml, the orientation of the orbital within the subshell. This can be between -l and l (max 2 e-)
10
Q
Spin quantum number
A
ms, used to distinguish between 2 electrons in an orbital, designated either +1/2 and -1/2. Electrons with different orbitals, but the same spin, are parallel
11
Q
n+l rule
A
Used to find what subshell fills first
12
Q
Hund’s rule
A
E- fill each subshell before they pair with each other in those subshells (half filled subshells)
13
Q
Paramagnetic
A
Elements with unpaired e- (the spin of the unpaired is parallel to each other), IN alignment with magnetic field causing weak attraction
14
Q
Diamagnetic
A
Elements with paired e-, repelled by magnetic field, Ex/ wood or plastic
15
Q
Heisenberg uncertainty principle
A
Inability to know the position and momentum of a single electron simultaneously. The position is given by the radius of the orbit
16
Q
Pauli exclusion principle
A
No two electrons can have the same four quantum numbers or else they would be occupying the same space
17
Q
Aufbau principle
A
Electrons will fill lower energy orbitals first before filling higher energy orbitals
*from the German Aufbauprinzip (building-up principle)
18
Q
Mass defect
A
Δm = nucleus mass - (mass of protons + mass of neutrons)
19
Q
Effective nuclear charge (Zeff)
A
strength of the electrostatic attraction between valence e- and the nucleus (protons). Trend?
20
Q
Shielding
A
Valence e- are increasingly separated from the nucleus by inner shells and the outermost e- are held less tightly. This is the effect of the trend for principal quantum number.
21
Q
Atomic radius
A
The distance between center of the nucleus and the outermost electron. Decreases from left to right and increases top to bottom (unique trend)
22
Q
Ionic radius
A
Half the distance between 2 ions that briefly make contact with each other. Cations will have a larger ionic radius than atomic. Anions will have a smaller ionic radius than atomic
23
Q
Ionization energy (IE)
A
energy required to remove an e- from the outer shell of an atom. Endothermic. Completing or disrupting the shell/subshell makes the strength low or high. Trend?
24
Q
Electron affinity
A
energy released when an e- is gained. Exothermic. Compare trend to Zeff and atomic radius.
25
Electronegativity
the attractive force of an element on an e- in a chemical bond. Rlship to IE?
26
Noble gases
No tendency to gain or lose e-. No measurable electronegativities. Extremely low boiling points. Exist as gas at room temp. London dispersion forces
27
Transition metals
Conductive and free moving valence e-. Ability to have multiple positive oxidation states from losing e- in their s or d subshells. Form complexes w/ water (hydration) and non-metals.
28
Metalloids
Can act as both metals and non metals depending on what they are bound to. Semi-conductors.
(8), Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Polonium
29
Ionic bond
bond between an atom of low e- affinity and an atom of high e- affinity. It results from an attraction between opposite charges. Rlship to melting/boiling points?
30
Covalent bond
an e- pair is shared between 2 atoms. Bond energy and number of e- pairs have an direct rlship.
31
Dipole moment (p)
p = qd —> q=charge and d=displacement vector. p in Columbs/meter
32
Coordinate covalent bond
formed when a lone pair attacks an unhybridized p-orbital. Ex/ Lewis rxn
33
Formal charge (FC)
FC = (Ve-) - (# nonbonding) - (1/2 # bonding)
- the difference between the number of valence e- of an atom in a particular Lewis structure and the number of valance e- normally found in that same atom.
34
Bent (angle)
104.5 Ex/ H2O
35
Trigonal pyrimidal
107. Ex/ NH3
36
Tetrahedral (angle)
109.5. Ex/ CH4
37
Trigonal planar
120. Ex/ SO3, CH2O
38
Trigonal bipyrmidal
90, 120, 180. Ex/ PCl5
39
Octahedral
90, 180. Ex/ AlF6
40
London dispersion (van der Waals) forces
weak interactions between a bond and other e- clouds that come within proximity. Rlship to vapor pressure?
41
Dipole-dipole forces
Direct interaction between 2 magnetic poles. It is an intermolecular force present in solid and liquid phases but negligible in gases
42
Hydrogen bonds
Occurs in O, N, and F
43
Equivalent
The amount of acid or base required to produce or consume 1 mole of protons. Describes how many moles of something we are interested in.
44
Gram equivalent weight
the amount of a compound (measured in grams), that produces one equivalent
GEW = molar mass/n
45
Normality
measure of concentration given in equivalents/L, commonly used for H+ concentration in acids
46
Empirical formula
the simplest whole number ratio of the elements in a compound
47
Molecular formula
the exact # of atoms of each element in a compound
48
Percent composition
% comp = (mass of element in formula* / molar mass ) x 100
*Assume a 100g when there are unknowns
49
Combustion reaction
CH4 + 2 O2 —> CO2 + 2 H2O
50
Limiting reagent
the reactant that has less moles when converted to a product and is used up first
51
Excess reagent
The reactant that has more moles than the other reactant after being converted into a product
52
Percent yield
% yield = (Actual / Theoretical) x 100
53
Rate determining step
Slow step, the rate of the whole reaction is only as fast as this step
54
Arrhenius reaction
```
k = Aℇ^(-Ea/RT)
A = frequency factor (attempt frequency) which is a measure of how often molecules collide
```
55
Activation complex
Energy of the transition state
56
Exergonic reaction
a negative free energy change, spontaneous
57
Endergonic reaction
a positive free energy change, non-spontaneous
58
Factors effecting reaction rate
Concentration, temperature, partial pressure, medium, catalysts (increase forward and reverse rxn)
59
Zero order reaction
The rate of formation of product C is independent of changes of any of the reactants A and B. Only effected by temp.
Rate = k[A]⁰[B]⁰ = k
60
First order reaction
The rate is directly proportional to only 1 reactant. Doubling the concentration of that reactant results in doubling the rate of formation of the product
Rate = k[A]¹
61
Second order reaction
The rate is proportional to either the concentration of 2 reactants OR the square of 1 reactant (Ex/ quartering [A] results in halving the rxn rate)
Rate = k[A]¹[B]¹ or rate = k[A]²
62
Reaction quotient
indicates how far the reaction has proceeded towards equilibrium by giving the concentrations of reactants and products at a certain point. Rlship to Keq?
63
LeChatelier's Principle
If a stress is applied to one side of a system, the system shifts to balance that stress. Rlship to volume, pressure, heat (endo/exo)?
64
Kinetic product
formed from a reactant at low temperatures with a small heat transfer. Less stable
65
Thermodynamic product
formed at high temperatures with a large heat transfer. More stable
66
Kd
Equilibrium constant of the dissociation reaction between receptor and ligand. A smaller Kd = a higher concentration of reactants (receptor-ligand complex) or a higher affinity between receptor and ligand
67
Ksp
Equilibrium constant of the solubility of ionic molecules in aqueous solutions. Ksp is temperature dependent so increasing temp will increase solubility (for non-gas solutes. In this case, increasing temp will decrease solubility for non-gas solutes. Ksp is higher for gas solutes at high pressure)
68
Normal body temperature
37℃ or 310 K
69
Liters in a mol (volume state function)
22.4 L = 1 mol
70
Standard conditions
25℃ (298 K), 1 atm, 1 M concentrations. Used for kinetics, equilibrium, and thermodynamics problems
71
Standard temperature and pressure
0℃ (273 K) and 1 atm. Used for ideal gas calculations. 1 mol of any gas at STP = 22.4 L
72
Triple point
temperature and pressure of all 3 phases are in equilibrium
73
Critical point
where the phase boundary between the liquid and gas phase terminates and there is no distinction between phases
74
Supercritical fluid
any substance at a temperature and pressure above its critical point. Distinct liquid and gas phases do not exist
75
Temperature
average kinetic energy of particles of a substance
76
Heat
q, transfer of energy from one substance to another as a result of their differences in temperature
77
Enthalpy
ΔH, heat under constant pressure. Rlship to endo/exo rxns?
78
Heating curve equations
q= mL when no phase change
| q=mcΔT during a phase change
79
Hess' Law
Enthalpy changes of reactions are additive
ΔHᵒrxn = ΔHᵒ(products) - ΔHº(reactants)
*Make sure to balance the equation!
80
Entropy
measure of spontaneous dispersal of energy at specific temperatures.
ΔS = Qrev/T with units of J/mol·K
81
Gibbs free energy (5 eqs)
```
ΔG = ΔH - TΔS
ΔGrxn = ΔG°rxn + RT ln(Q)
ΔGrxn = RT ln(Q/K𝘦𝑞)
ΔG°rxn = -RT· ln(K𝘦𝑞)
ΔG° = −nF E°cell
```
82
Ideal gas law
PV=nRT
R = 8.21 x 10⁻² L·atm/mol·K or 8.314 J/mol·K
Individual molecular volume and intermolecular forces are negligible
83
Avagadro's principle
All gases at a constant temperature and pressure occupy volumes that are directly proportional to the number of moles present
k = n/V or n₁/V₁ = n₂/V₂
84
Boyle's Law
P₁V₁ = P₂V₂
85
Charles' Law
V₁/T₁ = V₂/T₂
86
Gay-Lussac's Law
P₁/T₁ = P₂T₂
87
Dalton's Law of partial pressure
When 2 or more gases that do no interact are found in 1 vessel, the pressure in the vessel is the combined partial pressures that each gas exerts individually.
Ptotal = PA + PB + PC… where PA = XA(Ptotal)
88
Vapor pressure
pressure exerted by evaporated particles above the surface of a liquid
89
Henry's Law
solubility and vapor pressure have an direct relationship
| [A]₁/P₁ = [A]₂/P₂ = kH (Henry's constant)
90
Kinetic molecular theory
The average kinetic energy of a gas particle is proportional to the absolute temperature of the gas.
KE = 1/2 mv² = 3/2 kBT
91
Root mean squared
defines the average speed of a molecule to determine the average kinetic energy per particle and calculate the speed that corresponds.
𝜇RMS = √3RT/M
92
Graham's Law
Heavier gasses diffuse or effuse slower than lighter ones because of their lower average speed
r₁/r₂ = √M₂/M₁
93
Solvation
the electrostatic interaction between solvent and solute molecules, dissolution
94
Ideal solution
when the strength of the original interactions equals the strength of the interactions after solvation. The overall enthalpy change is 0
95
Solubility
the maximum amount of a substance that can be dissolved in a particular solvent at a given temp.
96
Saturated solution
the maximum amount of solute has been added to a soln
97
Precipitate
Formed when more solute is added after saturation. Transition metals that form precipitates have unfilled orbitals
98
Complex ion
a molecule that has a cation bond to at least one electron pair donor (ligand). Held together by coordinate covalent bonds
99
Chelation
where the central cation in a complex is bound to the same ligand in multiple places. They are usually large and used to sequester toxic metals
100
Percent composition by mass
(mass of solute) / (mass of solution) x 100
101
Mole fraction
Xa = (moles of A) / (total moles of species)
102
Molarity
M = moles of solute / liters of solution
103
Molality
m = moles of solute / kg of solvent
104
Normality
The number of equivalents (molecules) of interest per liter of solution. It is like the molarity of the molecules of interest
105
Solubility product constant
(Ksp) is the equilibrium constant for the solubility of a saturated solution of an ionic compound in an aqueous solution
Ksp = [Aⁿ⁺]ᵐ[Bᵐ⁻]ⁿ
106
Common ion effect
If a solute tries to dissolve in a solution that contains one of its ions, the molar solubility will decrease and the system will shift to the left and start to reform the solute
107
Strong electrolytes
dissolve completely into its ions
108
Weak electrolytes
do not completely dissolve into ions
109
Colligative properties
Freezing point depression, boiling point elevation, and osmotic pressure
110
Boiling point elevation
The boiling point can be raised when a non-volatile solute is dissolved in a solvent.
ΔTb = iKbm
111
Freezing point depression
The presence of solute particles can interfere with the formation of the lattice arrangement of a solid state so a greater amount of energy must be removed from the solution.
ΔTf = iKfm
112
Osmotic pressure
The pressure generated by solutions in which water is drawn into a that solution.
Π = iMRT
113
Miscibility
property when 2 substances mix in all proportions and form a homologous solution
114
Arrhenius acid
dissociates to form an excess of H⁺ in solution
115
Arrhenius base
dissociates to form an excess of OH⁻ in solution
116
Bronsted-Lowry acid
donates H⁺ ions
117
Bronsted-Lowry base
accepts H⁺ ions
118
Lewis acid
e⁻ pair acceptor
119
Lewis base
e⁻ pair donor
120
Amphoteric
species that can react like an acid in a basic environment or react like a base in an acidic environment
121
Amphiprotic
species that can can either gain or lose a proton
122
Equilibrium constant for ionization of water
Kw = [H₃O⁺][OH⁻] = 10⁻¹⁴ at 25℃
123
pH
measure of acidic concentration
| pH = -log[H+] or log 1/[H+]
124
pOH
measure of basic concentration
| pOH = -log[OH-] or log 1/[OH-]
125
Acid dissociation constant
HA + H₂O --> H₃O⁺ + A⁻
| Ka = [H₃O⁺][A⁻] / [HA]
126
Base dissociation constant
BOH --> B⁺ + OH⁻
| Kb = [B⁺][OH⁻] / [BOH]
127
Conjugate acid
acid that forms when a base gains a proton
128
Conjugate base
base that forms when an acid donates a proton
129
Induction
electronegative elements will increase acid strength (cause them to dissociate) by pulling the electron density away from the acid and increasing the ability for the proton on the acid to be lost.
130
Neutralization
when acids and bases react with each other to form a salt and water
131
Hydrolysis
when salt ions react with water to form an acid and base
132
Polyvalent
a mole of an acid or base can liberate more than 1 acid or base equivalent (H+ or OH-)
133
Titration
Performed by adding volumes titrant to a known volume of titrand until the equivalence point. Used to find the pKa/pKb or the unknown concentration of a substance
134
Equivalence point
Point in the titration when there are equal equivalents of acid and base in the solution
V1N1 = V2N2
135
Indicator
fully protonated or deprotonated weak acid/ bases (HIn) whose conjugate base/acid (I-) are a different color. It is best to use an indicator with a pKa as close as possible to the pH of the titration equivalence point
136
Buffer
molecules that help maintain pH in a normal range. They can contain mixtures of weak acid/conjugate base+cation or weak base/conjugate acid+anion
137
Henderson—Hasselbalch equation
Used to estimate the pH (or pOH) or a buffer solution.
| pH = pKa + log( [A⁻]/[HA] )
138
Oxidation
Loss of e-, decrease in the number of bonds to H and increase in the number of bonds to C, N, O, or halides, increase in the # of bonds to oxygen
139
Oxidizing agent
species that causes another species to be oxidized (it gets reduced)
140
Reduction
gain of e-, increasing the number of bonds to H and decreasing the number of bonds to C, N, O, or halides
141
Reducing agent
causes another agent to be reduced (it gets oxidized)
142
Galvanic (voltaic) cells
Convert chemical energy into electrical. The reactions are spontaneous, they have a decreasing ΔG, and a pos emf
143
Salt bridge
prevents the excess (+) and (-) charge build up at the anode and cathode when the cations and anions flow into their respective places. This minimizes potential difference. Contains inert electrolytes
144
Electrolytic cells
Convert electric energy into chemical. The rxns are non-spontaneous, ΔG is increasing, and there is a neg emf
145
Electrolysis
type of oxi-red reaction driven by an external voltage source where chemical compounds are decomposed
146
Electrode charge destinations
In a galvanic cell, the anode is negative and the cathode is positive.
In an electrolytic cell, the anode is positive and the cathode is negative
147
Faraday's constant
1 F = 96,485 C or 10⁵ C/mol e-
148
Electrodeposition equation
Tells us the # of moles of element being deposited on a plate
Mol M = It / nF --> I is current and n is the # of e- equivalents for a specific metal ion
149
Concentration cells
The electrodes are identical (same potentials). So that means the current is generated from the concentration gradient of ions in each half cell.
150
Standard state cell potential (electromotive force) equation
The sum of the standard state potentials of the half reactions.
E°cell = E°red (cathode - anode) or E°red(reduction - oxidation)
151
Nernst equation
E cell = E°cell - 0.0592/n (logQ)
Q = [X] anode / [Y] cathode
152
Autoionization of water
Water is amphoteric so it can react with itself to form hydronium and hydroxide ions
H2O + H2O --> H3O+ + OH-
153
Strong acid
Acid with a Ka > 1 or a pKa < 0. The larger the Ka and the smaller the pKa, the stronger the acid
154
Strong base
Base with a Kb < 1 or a pKb > 0. The smaller the Kb and the larger the pKb, the stronger the base
155
Gas constant R values
8.3 J/mol·K and 0.08 L·atm/K·mol
156
Titrant
the acid or base that is added to the substance of unknown concentration
157
Half equivalence point
The point where 1/2 of the acid has been neutralized by the base (the [acid] = [conjugate base]). It shows where the solution is most well buffered
158
Endpoint
The point where the indicator changes color. This is not the same as the equivalence point
159
Cell potential
E, also called the electromotive force (emf). The potential difference between 2 terminals when not connected. Drop in emf increases as the current increases.
160
Potential difference
The difference in electric potential between 2 points in a circuit
161
Reduction potential
a measure of the tendency of a chemical species to gain electrons from or lose electrons to an electrode and thereby be reduced or oxidized respectively
162
Amphiprotic vs. amphoteric vs. amphipathic
Species that can gain or lose a proton vs. a species that can react like an acid or base vs. a species that has hydrophilic and hydrophobic regions
163
Suspension
A heterogeneous mixture of larger particles that are visible and will settle out after the mixture stands. NOT a solution. Ex/ tomato juice
164
Solution
Homogeneous mixture of two substances (solute and solvent) that is uniform throughout in its appearance and composition
165
Colloid
Homogeneous mixture where a noncrystalline solute of large particles or microscopic particles is dispersed throughout the solvent. The particles do not settle. This is a solution. Ex/ fog, gel
166
Potentiometric titration
Measures the changes in voltage during a titration, since changes in concentration can affect voltage. A voltmeter is used
