CST study cards Flashcards
(222 cards)
1
Q
Oxidation number for alkali metal
A
+1
2
Q
Oxidation number for alkaline earth metal
A
+2
3
Q
Oxidation number for group IIIA metals
A
+3
4
Q
Oxidation number for hydrogen
A
+1
5
Q
Oxidation number for fluorine
A
-1
6
Q
Oxidation number for oxygen
A
-2
7
Q
Oxidation number for halogens
A
-1
8
Q
Oxidation number for Group VIA nonmetals
A
-2
9
Q
Very active metals
A
Li, Na, K, Rb, Cs, Ca, Sr, Ba
React with H2O to produce H2
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
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
12
Q
Activity series for halogens
A
F2 > Cl2 > Br2 > I2
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
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
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
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.
17
Q
Mercury (I)
formula and oxidation state
A
Hg_2^2+
18
Q
Ammonium
formula and oxidation state
A
NH_4^+
19
Q
Nitrite
formula and oxidation state
A
NO_2^-
20
Q
Nitrate
formula and oxidation state
A
NO_3^-
21
Q
Sulfite
formula and oxidation state
A
SO_3^2-
22
Q
Sulfate
formula and oxidation state
A
SO_4^2-
23
Q
Hydrogen sulfate (bisulfate) formula and oxidation state
A
HSO_4^-
24
Q
Hydrogen phosphate
formula and oxidation state
A
HPO_4^2-
25
Dihydrogen phosphate
| formula and oxidation state
H_2PO_4^-
26
Thiocyanate
| formula and oxidation state
SCN^-
27
Carbonate
| formula and oxidation state
CO_3^2-
28
```
Hydrogen carbonate (bicarbonate)\
formula and oxidation state
```
HCO_3^-
29
Hypochlorite
| formula and oxidation state
ClO^-
30
Chlorite
| formula and oxidation state
ClO_2^-
31
Chlorate
| formula and oxidation state
ClO_3^-
32
Perchlorate
| formula and oxidation state
ClO_4^-
33
Acetate
| formula and oxidation state
C_2H_3O_2^- or CH_3COO^-
34
Permanganate
| formula and oxidation state
MnO_4^-
35
Dichromate
| formula and oxidation state
Cr_2O_7^2-
36
Chromate
| formula and oxidation state
CrO_4^2-
37
Peroxide
| formula and oxidation state
O_2^2-
38
Oxalate
| formula and oxidation state
C_2O_4^2-
39
Metalloid elements
```
B - boron
Si - silicon
Ge - germanium
As - arsenic
Sb - antimony
Te - tellurium
```
40
equation relating the speed of light to the wavelength and frequency
c = lambda * nu
| speed of light = wavelength * frequency
41
equation relating the energy to the frequency of light
```
E = h * nu
energy = planck's constant * frequency
```
42
definition of n (structure of the atom)
principal energy level
| n = 1 is closest to the nucleus
43
definition of l (structure of the atom)
```
azimuthal quantum number (sublevel number)
number of sublevels cannot be > n
value cannot be > n-1
0 = s
1 = p
2 = d
3 = f
```
44
definition of m_l (structure of the atom)
magnetic quantum number (orbital number)
number of orbitals = 2l + 1
values = -l to l
indicates shape and orientation of orbital
45
Allotrope
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)
46
atomic radius trends in the periodic table
left to right - decreasing radius due to effective nuclear charge
top to bottom - increasing radius due to increased energy levels
47
effective nuclear charge
| definition and trend in the periodic table
Total nuclear charge - non valence electrons
| increases from left to right across a period
48
first ionization energy trends in the periodic table
usually decreases from top to bottom in a group
| usually increases left to right in a period
49
binding energy equation
BE = energy of incoming photon - energy of emitted photoelectron
50
electron affinity
| definition and trend in the periodic table
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)
51
electronegativity
| definition and trend in the periodic table
describes the attraction of electrons by individual atoms
| increases diagonally from bottom left to top right (F has the highest, Fr has the lowest)
52
combustion reaction definition
organic compounds that react with oxygen to form CO_2 and water
53
single-replacement reaction definition
element reacts with a compound to form a different element and a new compound
54
double-replacement reaction definition
two compounds react and the cation in one compound replaces the cation in the second compound
55
neutralization reaction definition
double replacement reaction in which one compound is an acid and one is a base
56
synthesis reaction definition
two or more elements react to form a compound
57
formation reaction defintion
synthesis reaction with the product having a coefficient of 1
58
addition reaction definition
a simple molecule or an element is added to another molecule to form a new molecule
59
decomposition reaction definition
a large molecule decomposes into its elements or into smaller molecules
60
thiosulfate
| formula and oxidation state
S_2O_3^2-
61
soluble compounds based on cations
sodium
potassium alkali metals
ammonium
62
soluble compounds based on anions
nitrate (NO_3^-)
63
chemical driving forces for double-replacement reactions
1. formation of water
2. formation of a precipitate
3. formation of a non-ionic (covalent) compound such as organic acids or gases
64
formal charge definition
formal charge = number of valence e- - [number of non-bonding e- + 1/2 number of bonding e-]
65
dipole moment equation
dipole moment = q * r
| = difference in charge * distance between the two nuclei
66
Delta electronegativity
| equation and meaning
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
67
bond order definition
bond order = total number of bonds for a given element/ # of atoms bonded to that element
68
bond strength calculation
bond energy is equal to bond strength
bond energy = h* nu
= Planck's constant * frequency of vibration
69
sigma bonds
2 s orbitals
1 s and 1p orbital
2 p orbitals (1st p overlap)
only one per covalent bond
70
pi bonds
after 1 sigma bond is formed, subsequent bonds are pi bonds
| 2 p orbitals (sideways overlaps)
71
hybrid orbital description
combines sigma and p orbitals
| sp3 - 1s and 3p orbitals combined to form identical bonds, tetrahedron
72
Boyle’s Law
```
P_1V_1 = P_2V_2
PV = constant
```
73
Charles's Law
V_1/T_1 = V_2/T_2
absolute zero = x-intercept of this curve
V/T = constant
74
Guy-Lussac's Law
```
P_1/T_1 = P_2/T_2
P/T = constant
```
75
Avogadro's Principle
```
n_1/V_1 = n_2/V_2
n/V = constant
```
76
Kinetic Molecular Theory
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.
77
Pressure definition
P = F/A
78
Graham's Law of effusion
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
79
Ideal gas law
PV = nRT
| applies at low pressures and high temperatures (not near where gases condense)
80
Real gases
cooled and/or compressed
| distance between particles decreased dramatically
81
Dalton's law of partial pressures
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
82
London dispersion forces
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
83
dipole-dipole forces
attraction between the partial positive end of one dipole and the partial negative end of another dipolar molecule
84
hydrogen bonding
very strong dipole-dipole attractive forces observed exclusively in compounds that have an F, O, or N bonded directly to a hydrogen atom.
85
Strong electrolytes
HCl
HBr
HI
86
Weak acids
```
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)
```
87
Weak bases
related to ammonia
88
reaction quotient
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.
89
Q = K_c
reaction is at equilibrium
90
Q does not change over time
Q = K_c, reaction is at equilibrium
91
Q < K_c
reaction will move in the forward direction to reach equilibrium
92
Q > K_c
reaction will move in the reverse direction to reach equilibrium
93
K_p
equilibrium constant for gas-phase reactions (usually only reactions that take place entirely in the gas phase are written this way)
94
equation for relationship between K_p and K_c
K_p = K_c(RT)^(delta n_g)
95
rule of thumb for solubility of a salt in water
0.1M
96
K_sp
solubility product constant
| equilibrium constant for the dissolution of a salt in water
97
common ion effect
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
K_a
acid ionization constant
equilibrium constant that describes the ionization of the acid in water
can be used to determine the pH
99
K_b
base ionization constant
equilibrium constant that describes the ionization of the base in water
can be used to determine pH
100
K_f
formation constant
equilibrium constant for when metal ions react with anions to form complexes
K_f = 1/K_d
101
K_d
dissociation constant
equilibrium constant for when metal ion complexes dissociate into its parts
K_d = 1/K_f
102
Increase temperature for exothermic reactions
favors reactants, decreases K
103
Increase temperature for endothermic reactions
favors products, increases K
104
Decrease temperature for exothermic reactions
favors products, increases K
105
Decrease temperature for endothermic reactions
favors reactants, decreases K
106
Arrhenius equation
k = Ae^(-E_a/RT)
107
zero-order reaction
rate = k
| plot of concentration of reactant vs time is a straight line with the slope = -rate
108
first-order reaction
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
second-order reaction
rate = k[A]^2
rate = k[A][B]
plot of 1/[A] vs time is a straight line with slope = k
110
collision theory
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 = N*f_e*f_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
transition state theory
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
heat of reaction
delta H = PE_products - PE_reactants
endothermic --> delta H is positive
exothermic --> delta H is negative
113
catalyst effects on reactions
reduces the Ea, effectively increasing the number of collisions with sufficient energy, the reaction comes to equilibrium more quickly
114
equation for kinetic energy
KE = 1/2*m*v^2
115
equation for potential energy (gravitational)
PE = K_grav*(m_1*m_2/r)
116
equation for potential energy (electrostatic)
PE = K_elec*(q_1*q_2/r)
117
heat energy equation
q= C*m*delta_T
118
first law of thermodynamics
delta_E = q + w
energy change = heat + work
delta_E = PE_final - PE_initial
119
equation for work
```
work = force * distance moved
work = pressure * area * distance moved
work = pressure * volume changed
```
120
Hess's Law
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
delta G^o equation / meaning
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
delta G equation and meaning
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
Oxidation
loss of electrons
124
Reduction
gain of electrons
125
What happens at the cathode during electrolysis
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
what happens at the anode during electrolysis
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
cell diagram
written from anode to cathode
Electrodes written at the outside of the diagram
vertical lines represent phase changes
128
equation to calculate moles of an ion from electric current
moles of X = I*t/(n*F)
| moles of X = current * time/(mole of electrons * Faraday's constant)
129
equation for the standard cell potential as a function of Q
E_cell = E^o_cell - (RT/nF)*lnQ
at equilibrium, E_cell = 0
E^o_cell = (RT/nF)*lnK_eq
130
relationship between delta_G^o and E^o_cell
delta_G^o = -nFE^o_cell
131
Arrhenius theory
an acid adds hydrogen ions to a solution and a base adds hydroxide ions
132
Bronsted-Lowry theory
an acid is a proton donor and a base is a proton acceptor
133
strong acids
HCl, HBr, HI, HClO_4, HNO_3, H_2SO_4
134
weak acids
HF, H_2CO_3, H_3PO_4, H_3AsO_4, HClO_3, HClO_2, HClO
135
relationship between bond strength and acid strength
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
oxoacid
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
Li^+ flame color
deep red (crimson)
138
Na^+ flame color
yellow
139
K^+ flame color
pale violet
140
Ca^2+ flame color
Orange-red
141
Sr^2+ flame color
Red
142
Ba^2+ flame color
Yellow-green
143
Cu^2+ flame color
Blue-green
144
Cu^+ color in aqueous solution
Green
145
Cu^2+ color in aqueous solution
Blue
146
Fe^2+ color in aqueous solution
Yellow-green (depending on the anion)
147
Fe^3+ color in aqueous solution
Orange-red (depending on the anion)
148
Co^2+ color in aqueous solution
Pink
149
Cr^3+ color in aqueous solution
Violet (Cr(NO_3)_3) to Green (CrCl_3)
150
Ni^2+ color in aqueous solution
Green
151
Mn^2+ color in aqueous solution
Pink
152
MnO_4^- color in aqueous solution
Purple
153
CrO_4^2- color in aqueous solution
Yellow
154
Cr_2O_7^2- color in aqueous solution
Orange
155
FeSCN^2+ color in aqueous solution
Deep red
156
CoCl_4^2-
Blue
157
F_2 gas color
Pale yellow
158
Cl_2 gas color
green-yellow
159
Br_2 liquid color
deep red
160
I_2 color
metallic gray solid; violet gas
161
S_8 color
yellow solid
162
Cu color
red metallic solid
163
Au color
yellow metallic solid
164
NO_2 gas color
brown
165
Avogadro's number
6.022x10^23
166
Beer's law
A (absorbance) = a*b*c
where a = constant (absorptivity)
b = optical path length
c = concentration
167
Lewis acids and bases
acid: electron pair acceptor
base: electron pair donor
168
hydroxyl group
-OH
169
carbonyl group
-(C=O)-
170
carboxyl
-(C=O)-OH
171
amino group
-NH_2
172
amido group
-(C=O)-NH_2
173
alcohol compound
R-OH
174
Ether compound
R1-O-R2
| formed when reacting two alcohols
175
Aldehyde compound
R-(C=O)-H
176
Ketone compound
R1-(C=O)-R2
177
Organic acid
R-(C=O)-OH
178
Ester compound
R1-(C=O)-O-R2
| formed when reacting an organic acid and an alcohol
179
Amine compound
R-NH_2
180
Amide compound
R-(C=O)-NH_2
181
Haloalkane (Halide)
R-X where X = F, Br, Cl, or I
182
phosphate
PO_4^3-
183
cyanide
CN^-
184
thiosulfate
S_2O_3^2-
185
IUPAC aldehyde
-al
186
IUPAC ketone
-one
187
IUPAC ether
R1-oxy-R2
188
IUPAC alcohol
-ol
189
IUPAC amine
amino-
190
Haber process
N_2 + 3H_2 --> 2NH_3
191
First law of thermodynamics
energy cannot be created or destroyed
| E = q+w
192
Second law of thermodynamics
entropy of an isolated system always increases
193
Third law of thermodynamics
entropy of a system approaches a constant as the temperature approaches absolute zero
194
Bohr model of the atom
small dense nucleus surrounded by electrons in distinct energy levels around the nucleus, circular orbits (planetary model)
195
Rutherford model of the atom
Small dense positively charged nucleus with electrons orbiting in fixed, predictable paths and the atom is mostly empty space
Gold foil experiment
196
Schrodinger model of the atom
quantum mechanical model that predicts the likelihood of finding an atom in a certain position
197
molality
mole solute/kg solvent
198
transmutation
an element changing into another element through radioactive decay
199
alpha decay
particle ejects a helium nucleus, reduces atomic number by 2
200
beta decay
a high energy electron is ejected from an atom and a neutron is transformed into a proton, increases atomic number by 1
201
gamma decay
a high energy photon is ejected from the atom, no change in atomic number necessary
202
Coulomb's law
F_E = k*|(q_1*q_2)/r^2|
- the greater the difference in charge, the greater the force
- the smaller the distance between charges, the greater the force
203
lattice energy
increases as the magnitude of the charge increases
| decreases as the atomic radius increases
204
alloy properties
high electrical conductivity
high strength
high hardness
heat and corrosion resistant
205
delta_H^o (in terms of bonds)
= sum of bonds broken - sum of bonds formed
206
delta_G for phases changes
= 0
207
enthalpy of solution
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
relationship between atom size and bond energy
the smaller the size of the atom, the greater the energy of the bond it forms
209
expanded octets
molecules that have d subshells available can have more than 8 valance electrons more never more than 12
210
solubility rules
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
Henderson-Hasselbalch equation
pH = pK_a + log_10([A-]/[HA])
212
Oxides reacting with water trend
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
amino acid compound
an organic acid with one or more amino group
| building blocks of proteins
214
general formula alkanes
C_nH_2n+2
215
general formula for alkenes
C_nH_2n
216
general formula for alkynes
C_nH_2n-2
217
general formula for aromatic hydrocarbons
C_nH_2n-6
218
primary alcohol
zero or one carbon atom is bonded to the carbon with the -OH attached
219
secondary alcohol
two carbon atoms are bonded to the carbon with the -OH attached
220
tertiary alcohol
three carbon atoms are bonded to the carbon with the -OH attached
221
dihydroxy alcohol
contains two hydroxy groups
| 1, 2 ethanediol is antifreeze (or ethylene glycol)
222
trihydroxy alcohol
contains three hydroxy groups
| 1, 2, 3, propanetriol (or glycerol)
