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Chemistry Flashcards

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1
Q

Atomic number

A

of protons, Z

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2
Q

Atomic mass (mass number)

A

of protons + neutrons, A

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3
Q

Isotopes

A

Structures of the same element that differ in the number of neutrons

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4
Q

Atomic weight

A

Weighted average of all isotopes

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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

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6
Q

Avagadro’s number

A

6.6 x 10^23 molecules/mol

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7
Q

Principal quantum number

A

n, it describes the size of the orbital (energy level). The max # of e- is 2n^2

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8
Q

Azimuthal quantum number

A

l, the shape and # of subshells. Subshells 0-3 are designated spdf

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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-)

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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

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11
Q

n+l rule

A

Used to find what subshell fills first

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12
Q

Hund’s rule

A

E- fill each subshell before they pair with each other in those subshells (half filled subshells)

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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

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14
Q

Diamagnetic

A

Elements with paired e-, repelled by magnetic field, Ex/ wood or plastic

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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

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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

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17
Q

Aufbau principle

A

Electrons will fill lower energy orbitals first before filling higher energy orbitals

*from the German Aufbauprinzip (building-up principle)

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18
Q

Mass defect

A

Δm = nucleus mass - (mass of protons + mass of neutrons)

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19
Q

Effective nuclear charge (Zeff)

A

strength of the electrostatic attraction between valence e- and the nucleus (protons). Trend?

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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.

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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)

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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

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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?

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24
Q

Electron affinity

A

energy released when an e- is gained. Exothermic. Compare trend to Zeff and atomic radius.

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25
Q

Electronegativity

A

the attractive force of an element on an e- in a chemical bond. Rlship to IE?

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26
Q

Noble gases

A

No tendency to gain or lose e-. No measurable electronegativities. Extremely low boiling points. Exist as gas at room temp. London dispersion forces

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27
Q

Transition metals

A

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.

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28
Q

Metalloids

A

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

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29
Q

Ionic bond

A

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?

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30
Q

Covalent bond

A

an e- pair is shared between 2 atoms. Bond energy and number of e- pairs have an direct rlship.

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31
Q

Dipole moment (p)

A

p = qd —> q=charge and d=displacement vector. p in Columbs/meter

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32
Q

Coordinate covalent bond

A

formed when a lone pair attacks an unhybridized p-orbital. Ex/ Lewis rxn

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33
Q

Formal charge (FC)

A

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.
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34
Q

Bent (angle)

A

104.5 Ex/ H2O

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35
Q

Trigonal pyrimidal

A
  1. Ex/ NH3
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36
Q

Tetrahedral (angle)

A

109.5. Ex/ CH4

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37
Q

Trigonal planar

A
  1. Ex/ SO3, CH2O
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38
Q

Trigonal bipyrmidal

A

90, 120, 180. Ex/ PCl5

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39
Q

Octahedral

A

90, 180. Ex/ AlF6

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40
Q

London dispersion (van der Waals) forces

A

weak interactions between a bond and other e- clouds that come within proximity. Rlship to vapor pressure?

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41
Q

Dipole-dipole forces

A

Direct interaction between 2 magnetic poles. It is an intermolecular force present in solid and liquid phases but negligible in gases

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42
Q

Hydrogen bonds

A

Occurs in O, N, and F

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43
Q

Equivalent

A

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.

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44
Q

Gram equivalent weight

A

the amount of a compound (measured in grams), that produces one equivalent
GEW = molar mass/n

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45
Q

Normality

A

measure of concentration given in equivalents/L, commonly used for H+ concentration in acids

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46
Q

Empirical formula

A

the simplest whole number ratio of the elements in a compound

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47
Q

Molecular formula

A

the exact # of atoms of each element in a compound

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48
Q

Percent composition

A

% comp = (mass of element in formula* / molar mass ) x 100

*Assume a 100g when there are unknowns

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49
Q

Combustion reaction

A

CH4 + 2 O2 —> CO2 + 2 H2O

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50
Q

Limiting reagent

A

the reactant that has less moles when converted to a product and is used up first

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51
Q

Excess reagent

A

The reactant that has more moles than the other reactant after being converted into a product

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52
Q

Percent yield

A

% yield = (Actual / Theoretical) x 100

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53
Q

Rate determining step

A

Slow step, the rate of the whole reaction is only as fast as this step

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54
Q

Arrhenius reaction

A 
k = Aℇ^(-Ea/RT)
A = frequency factor (attempt frequency) which is a measure of how often molecules collide
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55
Q

Activation complex

A

Energy of the transition state

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56
Q

Exergonic reaction

A

a negative free energy change, spontaneous

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57
Q

Endergonic reaction

A

a positive free energy change, non-spontaneous

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58
Q

Factors effecting reaction rate

A

Concentration, temperature, partial pressure, medium, catalysts (increase forward and reverse rxn)

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59
Q

Zero order reaction

A

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

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60
Q

First order reaction

A

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

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61
Q

Second order reaction

A

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

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62
Q

Reaction quotient

A

indicates how far the reaction has proceeded towards equilibrium by giving the concentrations of reactants and products at a certain point. Rlship to Keq?

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63
Q

LeChatelier’s Principle

A

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)?

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64
Q

Kinetic product

A

formed from a reactant at low temperatures with a small heat transfer. Less stable

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65
Q

Thermodynamic product

A

formed at high temperatures with a large heat transfer. More stable

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66
Q

Kd

A

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

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67
Q

Ksp

A

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
Q

Normal body temperature

A

37℃ or 310 K

69
Q

Liters in a mol (volume state function)

A

22.4 L = 1 mol

70
Q

Standard conditions

A

25℃ (298 K), 1 atm, 1 M concentrations. Used for kinetics, equilibrium, and thermodynamics problems

71
Q

Standard temperature and pressure

A

0℃ (273 K) and 1 atm. Used for ideal gas calculations. 1 mol of any gas at STP = 22.4 L

72
Q

Triple point

A

temperature and pressure of all 3 phases are in equilibrium

73
Q

Critical point

A

where the phase boundary between the liquid and gas phase terminates and there is no distinction between phases

74
Q

Supercritical fluid

A

any substance at a temperature and pressure above its critical point. Distinct liquid and gas phases do not exist

75
Q

Temperature

A

average kinetic energy of particles of a substance

76
Q

Heat

A

q, transfer of energy from one substance to another as a result of their differences in temperature

77
Q

Enthalpy

A

ΔH, heat under constant pressure. Rlship to endo/exo rxns?

78
Q

Heating curve equations

A

q= mL when no phase change

q=mcΔT during a phase change

79
Q

Hess’ Law

A

Enthalpy changes of reactions are additive
ΔHᵒrxn = ΔHᵒ(products) - ΔHº(reactants)
*Make sure to balance the equation!

80
Q

Entropy

A

measure of spontaneous dispersal of energy at specific temperatures.
ΔS = Qrev/T with units of J/mol·K

81
Q

Gibbs free energy (5 eqs)

A 
ΔG = ΔH - TΔS
ΔGrxn = ΔG°rxn + RT ln(Q)
ΔGrxn = RT ln(Q/K𝘦𝑞)
ΔG°rxn = -RT· ln(K𝘦𝑞)
ΔG° = −nF E°cell
82
Q

Ideal gas law

A

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
Q

Avagadro’s principle

A

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
Q

Boyle’s Law

A

P₁V₁ = P₂V₂

85
Q

Charles’ Law

A

V₁/T₁ = V₂/T₂

86
Q

Gay-Lussac’s Law

A

P₁/T₁ = P₂T₂

87
Q

Dalton’s Law of partial pressure

A

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
Q

Vapor pressure

A

pressure exerted by evaporated particles above the surface of a liquid

89
Q

Henry’s Law

A

solubility and vapor pressure have an direct relationship

[A]₁/P₁ = [A]₂/P₂ = kH (Henry’s constant)

90
Q

Kinetic molecular theory

A

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
Q

Root mean squared

A

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
Q

Graham’s Law

A

Heavier gasses diffuse or effuse slower than lighter ones because of their lower average speed
r₁/r₂ = √M₂/M₁

93
Q

Solvation

A

the electrostatic interaction between solvent and solute molecules, dissolution

94
Q

Ideal solution

A

when the strength of the original interactions equals the strength of the interactions after solvation. The overall enthalpy change is 0

95
Q

Solubility

A

the maximum amount of a substance that can be dissolved in a particular solvent at a given temp.

96
Q

Saturated solution

A

the maximum amount of solute has been added to a soln

97
Q

Precipitate

A

Formed when more solute is added after saturation. Transition metals that form precipitates have unfilled orbitals

98
Q

Complex ion

A

a molecule that has a cation bond to at least one electron pair donor (ligand). Held together by coordinate covalent bonds

99
Q

Chelation

A

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
Q

Percent composition by mass

A

(mass of solute) / (mass of solution) x 100

101
Q

Mole fraction

A

Xa = (moles of A) / (total moles of species)

102
Q

Molarity

A

M = moles of solute / liters of solution

103
Q

Molality

A

m = moles of solute / kg of solvent

104
Q

Normality

A

The number of equivalents (molecules) of interest per liter of solution. It is like the molarity of the molecules of interest

105
Q

Solubility product constant

A

(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
Q

Common ion effect

A

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
Q

Strong electrolytes

A

dissolve completely into its ions

108
Q

Weak electrolytes

A

do not completely dissolve into ions

109
Q

Colligative properties

A

Freezing point depression, boiling point elevation, and osmotic pressure

110
Q

Boiling point elevation

A

The boiling point can be raised when a non-volatile solute is dissolved in a solvent.
ΔTb = iKbm

111
Q

Freezing point depression

A

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
Q

Osmotic pressure

A

The pressure generated by solutions in which water is drawn into a that solution.
Π = iMRT

113
Q

Miscibility

A

property when 2 substances mix in all proportions and form a homologous solution

114
Q

Arrhenius acid

A

dissociates to form an excess of H⁺ in solution

115
Q

Arrhenius base

A

dissociates to form an excess of OH⁻ in solution

116
Q

Bronsted-Lowry acid

A

donates H⁺ ions

117
Q

Bronsted-Lowry base

A

accepts H⁺ ions

118
Q

Lewis acid

A

e⁻ pair acceptor

119
Q

Lewis base

A

e⁻ pair donor

120
Q

Amphoteric

A

species that can react like an acid in a basic environment or react like a base in an acidic environment

121
Q

Amphiprotic

A

species that can can either gain or lose a proton

122
Q

Equilibrium constant for ionization of water

A

Kw = [H₃O⁺][OH⁻] = 10⁻¹⁴ at 25℃

123
Q

pH

A

measure of acidic concentration

pH = -log[H+] or log 1/[H+]

124
Q

pOH

A

measure of basic concentration

pOH = -log[OH-] or log 1/[OH-]

125
Q

Acid dissociation constant

A

HA + H₂O –> H₃O⁺ + A⁻

Ka = [H₃O⁺][A⁻] / [HA]

126
Q

Base dissociation constant

A

BOH –> B⁺ + OH⁻

Kb = [B⁺][OH⁻] / [BOH]

127
Q

Conjugate acid

A

acid that forms when a base gains a proton

128
Q

Conjugate base

A

base that forms when an acid donates a proton

129
Q

Induction

A

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
Q

Neutralization

A

when acids and bases react with each other to form a salt and water

131
Q

Hydrolysis

A

when salt ions react with water to form an acid and base

132
Q

Polyvalent

A

a mole of an acid or base can liberate more than 1 acid or base equivalent (H+ or OH-)

133
Q

Titration

A

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
Q

Equivalence point

A

Point in the titration when there are equal equivalents of acid and base in the solution
V1N1 = V2N2

135
Q

Indicator

A

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
Q

Buffer

A

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
Q

Henderson—Hasselbalch equation

A

Used to estimate the pH (or pOH) or a buffer solution.

pH = pKa + log( [A⁻]/[HA] )

138
Q

Oxidation

A

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
Q

Oxidizing agent

A

species that causes another species to be oxidized (it gets reduced)

140
Q

Reduction

A

gain of e-, increasing the number of bonds to H and decreasing the number of bonds to C, N, O, or halides

141
Q

Reducing agent

A

causes another agent to be reduced (it gets oxidized)

142
Q

Galvanic (voltaic) cells

A

Convert chemical energy into electrical. The reactions are spontaneous, they have a decreasing ΔG, and a pos emf

143
Q

Salt bridge

A

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
Q

Electrolytic cells

A

Convert electric energy into chemical. The rxns are non-spontaneous, ΔG is increasing, and there is a neg emf

145
Q

Electrolysis

A

type of oxi-red reaction driven by an external voltage source where chemical compounds are decomposed

146
Q

Electrode charge destinations

A

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
Q

Faraday’s constant

A

1 F = 96,485 C or 10⁵ C/mol e-

148
Q

Electrodeposition equation

A

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
Q

Concentration cells

A

The electrodes are identical (same potentials). So that means the current is generated from the concentration gradient of ions in each half cell.

150
Q

Standard state cell potential (electromotive force) equation

A

The sum of the standard state potentials of the half reactions.

E°cell = E°red (cathode - anode) or E°red(reduction - oxidation)

151
Q

Nernst equation

A

E cell = E°cell - 0.0592/n (logQ)

Q = [X] anode / [Y] cathode

152
Q

Autoionization of water

A

Water is amphoteric so it can react with itself to form hydronium and hydroxide ions
H2O + H2O –> H3O+ + OH-

153
Q

Strong acid

A

Acid with a Ka > 1 or a pKa < 0. The larger the Ka and the smaller the pKa, the stronger the acid

154
Q

Strong base

A

Base with a Kb < 1 or a pKb > 0. The smaller the Kb and the larger the pKb, the stronger the base

155
Q

Gas constant R values

A

8.3 J/mol·K and 0.08 L·atm/K·mol

156
Q

Titrant

A

the acid or base that is added to the substance of unknown concentration

157
Q

Half equivalence point

A

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
Q

Endpoint

A

The point where the indicator changes color. This is not the same as the equivalence point

159
Q

Cell potential

A

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
Q

Potential difference

A

The difference in electric potential between 2 points in a circuit

161
Q

Reduction potential

A

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
Q

Amphiprotic vs. amphoteric vs. amphipathic

A

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
Q

Suspension

A

A heterogeneous mixture of larger particles that are visible and will settle out after the mixture stands. NOT a solution. Ex/ tomato juice

164
Q

Solution

A

Homogeneous mixture of two substances (solute and solvent) that is uniform throughout in its appearance and composition

165
Q

Colloid

A

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
Q

Potentiometric titration

A

Measures the changes in voltage during a titration, since changes in concentration can affect voltage. A voltmeter is used

MCAT (7 decks)

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