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1
Q

Oxidation number for alkali metal

A

+1

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2
Q

Oxidation number for alkaline earth metal

A

+2

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3
Q

Oxidation number for group IIIA metals

A

+3

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4
Q

Oxidation number for hydrogen

A

+1

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5
Q

Oxidation number for fluorine

A

-1

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6
Q

Oxidation number for oxygen

A

-2

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7
Q

Oxidation number for halogens

A

-1

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8
Q

Oxidation number for Group VIA nonmetals

A

-2

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9
Q

Very active metals

A

Li, Na, K, Rb, Cs, Ca, Sr, Ba

React with H2O to produce H2

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10
Q

Active metals

A

Mg, Zn, Pb, Ni, Al, Ti, Cr, Fe, Cd, Sn, Co

React with acids to form H2, but not with H2O

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11
Q

Inactive metals

A

Ag, Au, Cu, Pt

DO NOT form H2 with acids; may react with concentrated oxidizing acids HNO3 and H2SO4 or aqu regia

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12
Q

Activity series for halogens

A

F2 > Cl2 > Br2 > I2

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13
Q

galvanic or voltaic cell

A 
produces energy
spontaneous redox reaction that creates a flow of electrons
cathode (+) / reduction
anode (-) / oxidation
e- flows from anode to cathode
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14
Q

electrolytic cell

A 
requires energy
redox reaction is forced to occur by adding electric energy
cathode (-) / reduction
anode (+) / oxidation
e- flows from anode to cathode
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15
Q

Equation for standard reduction potentials

A

Ecell = Ecathode - Eanode
Ecell = Ereduction - Eoxidation
negative E_cell = reaction is not thermodynamically favored as written
positive E_cell = reaction is thermodynamically favored as written

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16
Q

When given standard reduction potentials for two half reactions, how can you determine which one is most likely to be reduced in the full reaction?

A

The half reaction with the more positive standard reduction potential is more likely to be reduced.

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17
Q

Mercury (I)

formula and oxidation state

A

Hg_2^2+

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18
Q

Ammonium

formula and oxidation state

A

NH_4^+

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19
Q

Nitrite

formula and oxidation state

A

NO_2^-

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20
Q

Nitrate

formula and oxidation state

A

NO_3^-

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21
Q

Sulfite

formula and oxidation state

A

SO_3^2-

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22
Q

Sulfate

formula and oxidation state

A

SO_4^2-

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23
Q 
Hydrogen sulfate (bisulfate)
formula and oxidation state
A

HSO_4^-

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24
Q

Hydrogen phosphate

formula and oxidation state

A

HPO_4^2-

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25
Q

Dihydrogen phosphate

formula and oxidation state

A

H_2PO_4^-

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26
Q

Thiocyanate

formula and oxidation state

A

SCN^-

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27
Q

Carbonate

formula and oxidation state

A

CO_3^2-

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28
Q 
Hydrogen carbonate (bicarbonate)\
formula and oxidation state
A

HCO_3^-

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29
Q

Hypochlorite

formula and oxidation state

A

ClO^-

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30
Q

Chlorite

formula and oxidation state

A

ClO_2^-

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31
Q

Chlorate

formula and oxidation state

A

ClO_3^-

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32
Q

Perchlorate

formula and oxidation state

A

ClO_4^-

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33
Q

Acetate

formula and oxidation state

A

C_2H_3O_2^- or CH_3COO^-

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34
Q

Permanganate

formula and oxidation state

A

MnO_4^-

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35
Q

Dichromate

formula and oxidation state

A

Cr_2O_7^2-

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36
Q

Chromate

formula and oxidation state

A

CrO_4^2-

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37
Q

Peroxide

formula and oxidation state

A

O_2^2-

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38
Q

Oxalate

formula and oxidation state

A

C_2O_4^2-

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39
Q

Metalloid elements

A 
B - boron
Si - silicon
Ge - germanium
As - arsenic
Sb - antimony
Te - tellurium
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40
Q

equation relating the speed of light to the wavelength and frequency

A

c = lambda * nu

speed of light = wavelength * frequency

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41
Q

equation relating the energy to the frequency of light

A 
E = h * nu
energy = planck's constant * frequency
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42
Q

definition of n (structure of the atom)

A

principal energy level

n = 1 is closest to the nucleus

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43
Q

definition of l (structure of the atom)

A 
azimuthal quantum number (sublevel number)
number of sublevels cannot be > n
value cannot be > n-1
0 = s
1 = p
2 = d
3 = f
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44
Q

definition of m_l (structure of the atom)

A

magnetic quantum number (orbital number)
number of orbitals = 2l + 1
values = -l to l
indicates shape and orientation of orbital

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45
Q

Allotrope

A

element that has two or more distinct sets of chemical and physical properties
examples:
O_2 and O_3
C: graphite, diamond, bukminsterfullerene (C_60)

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46
Q

atomic radius trends in the periodic table

A

left to right - decreasing radius due to effective nuclear charge
top to bottom - increasing radius due to increased energy levels

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47
Q

effective nuclear charge

definition and trend in the periodic table

A

Total nuclear charge - non valence electrons

increases from left to right across a period

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48
Q

first ionization energy trends in the periodic table

A

usually decreases from top to bottom in a group

usually increases left to right in a period

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49
Q

binding energy equation

A

BE = energy of incoming photon - energy of emitted photoelectron

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50
Q

electron affinity

definition and trend in the periodic table

A

energy change that results from adding a electron to an atom

increases diagonally from bottom left to top right (F has the highest, Fr has the lowest)

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51
Q

electronegativity

definition and trend in the periodic table

A

describes the attraction of electrons by individual atoms

increases diagonally from bottom left to top right (F has the highest, Fr has the lowest)

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52
Q

combustion reaction definition

A

organic compounds that react with oxygen to form CO_2 and water

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53
Q

single-replacement reaction definition

A

element reacts with a compound to form a different element and a new compound

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54
Q

double-replacement reaction definition

A

two compounds react and the cation in one compound replaces the cation in the second compound

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55
Q

neutralization reaction definition

A

double replacement reaction in which one compound is an acid and one is a base

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56
Q

synthesis reaction definition

A

two or more elements react to form a compound

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57
Q

formation reaction defintion

A

synthesis reaction with the product having a coefficient of 1

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58
Q

addition reaction definition

A

a simple molecule or an element is added to another molecule to form a new molecule

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59
Q

decomposition reaction definition

A

a large molecule decomposes into its elements or into smaller molecules

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60
Q

thiosulfate

formula and oxidation state

A

S_2O_3^2-

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61
Q

soluble compounds based on cations

A

sodium
potassium alkali metals
ammonium

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62
Q

soluble compounds based on anions

A

nitrate (NO_3^-)

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63
Q

chemical driving forces for double-replacement reactions

A
  1. formation of water
  2. formation of a precipitate
  3. formation of a non-ionic (covalent) compound such as organic acids or gases
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64
Q

formal charge definition

A

formal charge = number of valence e- - [number of non-bonding e- + 1/2 number of bonding e-]

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65
Q

dipole moment equation

A

dipole moment = q * r

= difference in charge * distance between the two nuclei

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66
Q

Delta electronegativity

equation and meaning

A

delta electronegativity = (atom with largest electronegativity) - (atom with smallest electronegativity)
the greater the delta EN, the more polar the bond
if delta EN = 0, the bond is non-polar
delta EN > 1.7, bond is ionic
delta EN < 1.7, bond is polar covalent

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67
Q

bond order definition

A

bond order = total number of bonds for a given element/ # of atoms bonded to that element

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68
Q

bond strength calculation

A

bond energy is equal to bond strength
bond energy = h* nu
= Planck’s constant * frequency of vibration

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69
Q

sigma bonds

A

2 s orbitals
1 s and 1p orbital
2 p orbitals (1st p overlap)
only one per covalent bond

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70
Q

pi bonds

A

after 1 sigma bond is formed, subsequent bonds are pi bonds

2 p orbitals (sideways overlaps)

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71
Q

hybrid orbital description

A

combines sigma and p orbitals

sp3 - 1s and 3p orbitals combined to form identical bonds, tetrahedron

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72
Q

Boyle’s Law

A 
P_1V_1 = P_2V_2
PV = constant
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73
Q

Charles’s Law

A

V_1/T_1 = V_2/T_2
absolute zero = x-intercept of this curve
V/T = constant

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74
Q

Guy-Lussac’s Law

A 
P_1/T_1 = P_2/T_2
P/T = constant
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75
Q

Avogadro’s Principle

A 
n_1/V_1 = n_2/V_2
n/V = constant
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76
Q

Kinetic Molecular Theory

A
  1. Gases consist of molecules or atoms in continuous motion.
  2. Collisions between these molecules and/or atoms in a gas are elastic.
  3. The volume occupied by the atoms and/or molecules in a gas are negligibly small.
  4. The attractive or replusive forces between the atoms and/or molecules in a gas are negligible.
  5. The average kinetic energy of a molecule or atom in a gas is directly proportional to the Kelvin temperature of the gas.
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77
Q

Pressure definition

A

P = F/A

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78
Q

Graham’s Law of effusion

A

sqrt (m1/m2) = v_rms2/v_rms1
square root of the mass of molecule 1 / the mass of molecule 2 = the rate of the diffusion of molecule 2 / the rater of the diffusion of molecule 1

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79
Q

Ideal gas law

A

PV = nRT

applies at low pressures and high temperatures (not near where gases condense)

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80
Q

Real gases

A

cooled and/or compressed

distance between particles decreased dramatically

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81
Q

Dalton’s law of partial pressures

A

when two gases are mixed together, the gas particles tend to act independently of each other
P_total = P_1 + P_1 + … were P stands for the partial pressure of each individual gas

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82
Q

London dispersion forces

A

dispersion forces / instantaneous dipoles / induced dipoles
weak attractive forces due to the momentary unequal distribution of electrons around an atom
the larger the molecule, the greater the London dispersion forces

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83
Q

dipole-dipole forces

A

attraction between the partial positive end of one dipole and the partial negative end of another dipolar molecule

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84
Q

hydrogen bonding

A

very strong dipole-dipole attractive forces observed exclusively in compounds that have an F, O, or N bonded directly to a hydrogen atom.

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85
Q

Strong electrolytes

A

HCl
HBr
HI

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86
Q

Weak acids

A 
ethanoic acid / acetic acid (HC_2H_3O_2 or CH_3COOH)
Methanoic acid (HCHO_2 or HCOOH)
Propanoic acid (HC_3H_5O_2 or CH_3CH_2COOH)
Benzoic acid (HC_7H_5O_2 or C_6H_5COOH)
hypochlorous acid (HClO or HOCl)
chlorous acid (HClO_2 or HOClO)
Chloric acid (HClO_3, or HOClO_2)
hydrosulfuric acid (H_2S)
hydrofluoric acid (HF)
phosphoric acid (H_3PO_4)
water (H_2O)
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87
Q

Weak bases

A

related to ammonia

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88
Q

reaction quotient

A

Q
Use the equilibrium expression to calculate Q, compare Q to K_c to determine if the reaction is at equilibrium and if the forward or reverse reaction is favored.

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89
Q

Q = K_c

A

reaction is at equilibrium

90
Q

Q does not change over time

A

Q = K_c, reaction is at equilibrium

91
Q

Q < K_c

A

reaction will move in the forward direction to reach equilibrium

92
Q

Q > K_c

A

reaction will move in the reverse direction to reach equilibrium

93
Q

K_p

A

equilibrium constant for gas-phase reactions (usually only reactions that take place entirely in the gas phase are written this way)

94
Q

equation for relationship between K_p and K_c

A

K_p = K_c(RT)^(delta n_g)

95
Q

rule of thumb for solubility of a salt in water

96
Q

K_sp

A

solubility product constant

equilibrium constant for the dissolution of a salt in water

97
Q

common ion effect

A

decrease in solubility of a compound when it is dissolved in a solution that contains an ion in common with the salt being dissolved.

98
Q

K_a

A

acid ionization constant
equilibrium constant that describes the ionization of the acid in water
can be used to determine the pH

99
Q

K_b

A

base ionization constant
equilibrium constant that describes the ionization of the base in water
can be used to determine pH

100
Q

K_f

A

formation constant
equilibrium constant for when metal ions react with anions to form complexes
K_f = 1/K_d

101
Q

K_d

A

dissociation constant
equilibrium constant for when metal ion complexes dissociate into its parts
K_d = 1/K_f

102
Q

Increase temperature for exothermic reactions

A

favors reactants, decreases K

103
Q

Increase temperature for endothermic reactions

A

favors products, increases K

104
Q

Decrease temperature for exothermic reactions

A

favors products, increases K

105
Q

Decrease temperature for endothermic reactions

A

favors reactants, decreases K

106
Q

Arrhenius equation

A

k = Ae^(-E_a/RT)

107
Q

zero-order reaction

A

rate = k

plot of concentration of reactant vs time is a straight line with the slope = -rate

108
Q

first-order reaction

A

rate = k[A]
plot of concentration of reactant vs time decays exponentially
plot of log[A] vs time is a straight line with a slope of -k

109
Q

second-order reaction

A

rate = k[A]^2
rate = k[A][B]
plot of 1/[A] vs time is a straight line with slope = k

110
Q

collision theory

A

molecules must collide in exactly the right way, if that happens, the molecules will stop and all the kinetic energy will be converted into potential energy
If the potential energy exceeds the activation energy, the reaction will happen

reaction rate = Nf_ef_o
N = number of collisions per second
f_e = fraction of collisions with the minimum energy
f_o = fraction of collisions with correct orientation

111
Q

transition state theory

A

as molecules get closer, their orbitals interact and distort each other, weakening the bonds in the molecule, enabling new bonds to form
when the bonds are half-broken and half-formed, this is the activated complex

112
Q

heat of reaction

A

delta H = PE_products - PE_reactants
endothermic –> delta H is positive
exothermic –> delta H is negative

113
Q

catalyst effects on reactions

A

reduces the Ea, effectively increasing the number of collisions with sufficient energy, the reaction comes to equilibrium more quickly

114
Q

equation for kinetic energy

A

KE = 1/2mv^2

115
Q

equation for potential energy (gravitational)

A

PE = K_grav(m_1m_2/r)

116
Q

equation for potential energy (electrostatic)

A

PE = K_elec(q_1q_2/r)

117
Q

heat energy equation

A

q= Cmdelta_T

118
Q

first law of thermodynamics

A

delta_E = q + w
energy change = heat + work

delta_E = PE_final - PE_initial

119
Q

equation for work

A 
work = force * distance moved
work = pressure * area * distance moved
work = pressure * volume changed
120
Q

Hess’s Law

A
  1. if the coefficients of a chemical reaction are multiplied by a constant, the delta H^o_react is multiplied by the same constant
  2. If two or more equations are added together to obtain an overall reaction, the heats of these equations are also added to give the heat of the overall reaction.
121
Q

delta G^o equation / meaning

A

delta_G^o = delta_H^o - T*delta_S^o
negative delta_G^o = spontaneous / thermodynamically favored
postive delta_G^o = non-spontaneous / not thermodynamically unfavored

122
Q

delta G equation and meaning

A

delta_G = delta_G^o + RT lnQ
when Q = 1, delta_G = delta_G^o
delta_G >0 = reaction proceeds in reverse direction
delta_G <0 = reaction proceeds in the forward direction
delta_G = 0 = reaction is at equilibrium

at equilibrium,
delta_G^o = -RTlnK_eq

123
Q

Oxidation

A

loss of electrons

124
Q

Reduction

A

gain of electrons

125
Q

What happens at the cathode during electrolysis

A
  1. water (or H+) will be reduced to hydrogen gas if the other cations in the solution can be reduced to very active metals.
  2. If the metal ions can be reduced to inactive (or moderately active) metals, they will be reduced at the cathode instead of the water
126
Q

what happens at the anode during electrolysis

A
  1. if the anion is a polyatomic ion, it generally will not be oxidized (particularly sulfate, nitrate, and perchlorate)
  2. chloride, bromide, and iodide ions will be oxidized in aqueous solution
127
Q

cell diagram

A

written from anode to cathode
Electrodes written at the outside of the diagram
vertical lines represent phase changes

128
Q

equation to calculate moles of an ion from electric current

A

moles of X = It/(nF)

moles of X = current * time/(mole of electrons * Faraday’s constant)

129
Q

equation for the standard cell potential as a function of Q

A

E_cell = E^o_cell - (RT/nF)*lnQ
at equilibrium, E_cell = 0

E^o_cell = (RT/nF)*lnK_eq

130
Q

relationship between delta_G^o and E^o_cell

A

delta_G^o = -nFE^o_cell

131
Q

Arrhenius theory

A

an acid adds hydrogen ions to a solution and a base adds hydroxide ions

132
Q

Bronsted-Lowry theory

A

an acid is a proton donor and a base is a proton acceptor

133
Q

strong acids

A

HCl, HBr, HI, HClO_4, HNO_3, H_2SO_4

134
Q

weak acids

A

HF, H_2CO_3, H_3PO_4, H_3AsO_4, HClO_3, HClO_2, HClO

135
Q

relationship between bond strength and acid strength

A

the stronger the hydrogen is bonded to the acid, the weaker the acid
as electronegativity increases from left to right across a period, the weaker the bond gets, and the stronger the acid becomes
the strength of the acid increases from top to bottom in a group due to an increase in bond length (implies weaker bond)

136
Q

oxoacid

A

oxygen atoms bound to a central atom and hydrogen atoms are bound to the oxygen atoms
strength of the acid depends on the relative strength of the O-H bond.
O-H bond strength depends on
- number of O atoms (acid strength increases as the number of O atoms increases)
- the electronegativity of the central atom (acid strength increases as the electronegativity of the central atom increases)

137
Q

Li^+ flame color

A

deep red (crimson)

138
Q

Na^+ flame color

139
Q

K^+ flame color

A

pale violet

140
Q

Ca^2+ flame color

A

Orange-red

141
Q

Sr^2+ flame color

142
Q

Ba^2+ flame color

A

Yellow-green

143
Q

Cu^2+ flame color

A

Blue-green

144
Q

Cu^+ color in aqueous solution

145
Q

Cu^2+ color in aqueous solution

146
Q

Fe^2+ color in aqueous solution

A

Yellow-green (depending on the anion)

147
Q

Fe^3+ color in aqueous solution

A

Orange-red (depending on the anion)

148
Q

Co^2+ color in aqueous solution

149
Q

Cr^3+ color in aqueous solution

A

Violet (Cr(NO_3)_3) to Green (CrCl_3)

150
Q

Ni^2+ color in aqueous solution

151
Q

Mn^2+ color in aqueous solution

152
Q

MnO_4^- color in aqueous solution

153
Q

CrO_4^2- color in aqueous solution

154
Q

Cr_2O_7^2- color in aqueous solution

155
Q

FeSCN^2+ color in aqueous solution

156
Q

CoCl_4^2-

157
Q

F_2 gas color

A

Pale yellow

158
Q

Cl_2 gas color

A

green-yellow

159
Q

Br_2 liquid color

160
Q

I_2 color

A

metallic gray solid; violet gas

161
Q

S_8 color

A

yellow solid

162
Q

Cu color

A

red metallic solid

163
Q

Au color

A

yellow metallic solid

164
Q

NO_2 gas color

165
Q

Avogadro’s number

A

6.022x10^23

166
Q

Beer’s law

A

A (absorbance) = abc
where a = constant (absorptivity)
b = optical path length
c = concentration

167
Q

Lewis acids and bases

A

acid: electron pair acceptor
base: electron pair donor

168
Q

hydroxyl group

169
Q

carbonyl group

170
Q

carboxyl

171
Q

amino group

172
Q

amido group

A

-(C=O)-NH_2

173
Q

alcohol compound

174
Q

Ether compound

A

R1-O-R2

formed when reacting two alcohols

175
Q

Aldehyde compound

176
Q

Ketone compound

A

R1-(C=O)-R2

177
Q

Organic acid

A

R-(C=O)-OH

178
Q

Ester compound

A

R1-(C=O)-O-R2

formed when reacting an organic acid and an alcohol

179
Q

Amine compound

180
Q

Amide compound

A

R-(C=O)-NH_2

181
Q

Haloalkane (Halide)

A

R-X where X = F, Br, Cl, or I

182
Q

phosphate

183
Q

cyanide

184
Q

thiosulfate

185
Q

IUPAC aldehyde

186
Q

IUPAC ketone

187
Q

IUPAC ether

188
Q

IUPAC alcohol

189
Q

IUPAC amine

190
Q

Haber process

A

N_2 + 3H_2 –> 2NH_3

191
Q

First law of thermodynamics

A

energy cannot be created or destroyed

E = q+w

192
Q

Second law of thermodynamics

A

entropy of an isolated system always increases

193
Q

Third law of thermodynamics

A

entropy of a system approaches a constant as the temperature approaches absolute zero

194
Q

Bohr model of the atom

A

small dense nucleus surrounded by electrons in distinct energy levels around the nucleus, circular orbits (planetary model)

195
Q

Rutherford model of the atom

A

Small dense positively charged nucleus with electrons orbiting in fixed, predictable paths and the atom is mostly empty space
Gold foil experiment

196
Q

Schrodinger model of the atom

A

quantum mechanical model that predicts the likelihood of finding an atom in a certain position

197
Q

molality

A

mole solute/kg solvent

198
Q

transmutation

A

an element changing into another element through radioactive decay

199
Q

alpha decay

A

particle ejects a helium nucleus, reduces atomic number by 2

200
Q

beta decay

A

a high energy electron is ejected from an atom and a neutron is transformed into a proton, increases atomic number by 1

201
Q

gamma decay

A

a high energy photon is ejected from the atom, no change in atomic number necessary

202
Q

Coulomb’s law

A

F_E = k|(q_1q_2)/r^2|

  • the greater the difference in charge, the greater the force
  • the smaller the distance between charges, the greater the force
203
Q

lattice energy

A

increases as the magnitude of the charge increases

decreases as the atomic radius increases

204
Q

alloy properties

A

high electrical conductivity
high strength
high hardness
heat and corrosion resistant

205
Q

delta_H^o (in terms of bonds)

A

= sum of bonds broken - sum of bonds formed

206
Q

delta_G for phases changes

207
Q

enthalpy of solution

A

three parts

  1. energy needed to break the solute bonds (equal to lattice energy (always positive)
  2. energy needed to separate the water molecules (always positive)
  3. energy to create new associations between the solute and the dipoles of the water (always negative)

Step 2 and 3 together are called the hydration energy (always negative) and is a Coulombic energy

208
Q

relationship between atom size and bond energy

A

the smaller the size of the atom, the greater the energy of the bond it forms

209
Q

expanded octets

A

molecules that have d subshells available can have more than 8 valance electrons more never more than 12

210
Q

solubility rules

A

compounds with alkali metal cation (Na+, Li+, K+) or an ammonium cation (NH4+) are always soluble
compounds with a nitrate (NO3-) are always soluble

211
Q

Henderson-Hasselbalch equation

A

pH = pK_a + log_10([A-]/[HA])

212
Q

Oxides reacting with water trend

A

highest oxidation state for each atom
left hand side of the periodic table, the oxides are strongly basic
right hand side of the periodic table, the oxides are strongly acidic
oxides in the middle of the table are amphoteric (aluminum oxide as an example) having both acid and base properties

213
Q

amino acid compound

A

an organic acid with one or more amino group

building blocks of proteins

214
Q

general formula alkanes

215
Q

general formula for alkenes

216
Q

general formula for alkynes

217
Q

general formula for aromatic hydrocarbons

218
Q

primary alcohol

A

zero or one carbon atom is bonded to the carbon with the -OH attached

219
Q

secondary alcohol

A

two carbon atoms are bonded to the carbon with the -OH attached

220
Q

tertiary alcohol

A

three carbon atoms are bonded to the carbon with the -OH attached

221
Q

dihydroxy alcohol

A

contains two hydroxy groups

1, 2 ethanediol is antifreeze (or ethylene glycol)

222
Q

trihydroxy alcohol

A

contains three hydroxy groups

1, 2, 3, propanetriol (or glycerol)

CST Chemistry (1 decks)

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