Chem Flashcards
1
Q
How many cm^3 in a litre?
A
1000
2
Q
How many litres in a m^3?
A
1000
3
Q
What are the 7 rules for assigning oxidation states? (in order of precedence)
A
1) any element in its standard state is 0
2) the sum of the oxidation states in a molecule must equal the molecule’s charge
3) group 1 metals have +1, group 2 metals +2
4) fluorine is -1
5) H is +1 when bound to something more electronegative than C, -1 when bound to some less electronegative and 0 when bound to C
6) oxygen is -2
7) the rest of the halogens have -1 and atoms in the oxygen family have -2
4
Q
What are isotopes?
A
two of the same element that differ in their number of neutrons
5
Q
How are frequency and wavelength related to photon energy?
A
E = hf = hc/wavelength
directly proportional to frequency, indirectly proportional to wavelength
6
Q
Describe the Bohr model of the atom
A
electrons orbit the nucleus in circular paths
electrons with greater energy orbit at greater distances
electrons have quantized energy states
7
Q
What model helped explain that emission spectra are line spectra and not continuous spectra?
A
Bohr model
8
Q
What is a Bohr atom?
A
an atom that contains only one electron
9
Q
How would the energy emitted from an electron dropping from n=4 to n=3 compare to that of an electron dropping from n=3 to n=2?
A
n=4 to n=3 would have LESS energy released
i.e. the differences get smaller between higher levels
10
Q
What regions of the electromagnetic spectrum can be emitted when an electron transitions?
A
UV, visible, infrared
11
Q
What shape are s orbitals? How many electrons can they hold?
A
spherical
can hold 2 electrons
12
Q
What shape are p orbitals? How many electrons can they hold?
A
sort of like a dumbbell
can hold 6 electrons
13
Q
How many electrons can d and f subshells hold?
A
d can hold 10 electrons
f can hold 14 electrons
14
Q
What is the Aufbau principle?
A
electrons occupy the lowest energy orbitals available
15
Q
What is Hund’s rule?
A
electrons in the same subshell occupy available orbitals singly before pairing up
16
Q
What is the Pauli exclusion principle?
A
there can be no more than 2 electrons in any given orbital
17
Q
What does diamagnetic mean?
A
all electrons are spin-paired
there is no net magnetic field
atom will be repelled by an externally produced magnetic field
18
Q
What does paramagnetic mean?
A
electrons are NOT all spin-paired
atom will be attracted to an externally produced magnetic field
19
Q
What is a period? Group/family?
A
period = horizontal row across the periodic table group/family = vertical column
20
Q
What can happen that gives electron configurations different from what you would expect?
A
if you can have d5 or d10 i.e. a half or completely full shell an electron will be promoted from s2 to do so
21
Q
What does isoelectronic mean?
A
when two atoms have the same electron configuration
22
Q
How are electrons lost from orbitals?
A
always from the highest energy (highest n) first
ie 4s2 before 3d
23
Q
What is the excited state of an atom?
A
when electron configuration isn’t what is predicted but still follows the rules
24
Q
Which group are the alkali metals?
A
group 1
25
Which group are the alkaline earth metals?
group 2
26
Name the 7 metalloids
```
boron (B)
silicon (S)
Ge
As
Sb
Te
Po
```
27
What is nuclear shielding/the shielding effect?
each filled shell between the nucleus and the valence electrons shields the valence electrons from the full effect of the positively charged protons in the nucleus
28
What is the result of nuclear shielding?
effective nuclear charge
| valence electrons experience an effective reduction in the positive elementary charge
29
What are the trends for atomic radius?
decreases across a period
| increases down a group
30
How do ions vary in size?
cations
31
What is ionization energy?
the amount of energy necessary to remove the least tightly bound electron from an atom
32
What are the trends for ionization energy?
increases across a period
| increases up a group
33
How does first ionization energy compare to second?
second will always be greater
34
What is electron affinity?
the energy associated with the addition of an electron to an atom
if energy is released affinity is negative, if it is required affinity is positive
(negative is favoured)
35
What are the trends for electron affinity?
becomes more negative as you go across a period or up a group
(decreases to the right and up)
36
What is electronegativity?
a measure of an atom's ability to pull electrons to itself when it forms a covalent bond
37
What are the trends for electronegativity? Give specifically the order of the 9 most electronegative elements
increases across a period
increases up a group
F > O > N > Cl > Br > I > S > C ~ H
38
What are the trends for acidity?
increases across a period
| increases down a group
39
What are the 3 types of intramolecular forces?
covalent bonds
ionic bonds
metallic bonds
40
What is a coordinate covalent bond?
One atom donates both electrons in a bond
| ie hemoglobin and iron
41
What affects ionic bond strength? Which trend is most important?
1) increasing charge magnitude increases bond strength (most important)
2) decreasing ionic radius increases bind strength
42
What are some properties of metals? (with metallic bonds)
sea of delocalized valence electrons between nuclei
conduct heat and electricity
malleable
ductile
43
When making a Lewis dot diagram for a transition metal what is different about it?
you don't include valence electrons in the d orbitals
44
Which elements does the octet rule apply to?
2nd row
| B, C, N, O, F, Ne
45
What hybridization is H?
always just s because it can'r hybridize
46
What is the difference between orbital geometry and shape?
orbital geometry is determined by the # of electron groups
| shape is determined only by the bonds
47
What are the 3 intermolecular forces?
London Dispersion Forces (LDFs)
Dipole-dipole
H bonding
48
Rank the types of intermolecular forces from strongest to weakest
H bonding > dipole-dipole > LDFs
49
What are van der Waals forces?
includes H bonding, dipole-dipole and LDFs, but a lot of people call LDFs van der Waals
50
Describe London Dispersion Forces
present between all molecules
weakest intermolecular force
caused by instantaneous dipoles
increase with increasing molar mass
51
Describe dipole-dipole forces
only happen between polar molecules
| permanent dipole is attracted to another
52
What is needed for an H-bond?
H covalently bound to F, O or N
| Lone pair of electrons on F, O or N
53
What id homiletic bond cleavage?
one electron of the bond being broken goes to each fragment of the molecule (forms two radicals)
54
What is heterolytic bond cleavage?
both electrons of the bond being broken end up on the same atom
55
What is bond dissociation energy?
the energy required to break a bond homolytically
56
What is the teen between bond length and bond dissociation energy?
the higher the bond order (#of bonds between two atoms), the shorter and stronger the bond is
57
What are bond lengths measured in?
angstroms, A
| 1 x 10^-10m
58
How does s character related to the length of a bond?
the greater the s character in hybrid orbitals, the shorter the bond i.e. sp-sp bonds are shorter than sp-sp3
59
What is the relationship between intermolecular forces and melting point?
higher IMFs = higher melting point
| need to weaken IMFs in order to melt something
60
What is the relationship between intermolecular forces and boiling point
higher IMFs = higher boiling point
| need to break all IMFs in order for boiling to occur
61
What is vapour pressure? How is it related to intermolecular forces?
the pressure of a gas that has evaporated from liquid
| higher IMFs = LOW vapour pressure
62
What does volatile mean in terms of intermolecular forces?
low intermolecular forces, high vapour pressure
63
What is the relationship between viscosity and intermolecular forces?
high IMFs = high resistance = high viscosity
64
What is a network solid?
a solid in which atoms are connected in a lattice of covalent bonds
there are only intramolecular forces
very strong, tend to be very hard solids at room temp i.e. diamond
65
What is a metallic solid?
a metal
covalently bound lattice of nuclei and their inner shell electrons surrounded by a "sea" of electrons
vary widely in strength, but most are solid at room temp
66
What is a molecular solid?
crystal lattice of molecules held together by intermolecular interactions
often liquids or gases at room temp
67
What is the zeroth law of thermodynamics?
When systems are in thermal equilibrium, they have the same temperature
If two bodies of different temperatures are brought into contact with one another, heat will flow from the body with the higher temperature to the body with the lower temperature
68
What is the first law of thermodynamics?
energy cannot be created or destroyed, only transformed
| also, work can be put into a system to increase its overall energy (may or may not occur with a temperature change)
69
Describe enthalpy when bonds are broken and formed
when bonds are formed energy is released (H0), endothermic
70
What is delta H?
"heat of reaction"
71
What are the 3 ways to calculate delta H (heat of reaction)?
1) using heats of formation
(sum of n x delta H products) - (sum of n x delta H reactants)
2) Hess's law of heat summation
enthalpy is a state function, energy absorbed or released by overall reaction is the same as the sum of the individual steps
3) Summation of bond enthalpies (bond dissociation energies)
(sum BDEbroken) - (sum BDEformed)
72
What is is a heat of formation?
the change in energy associated with forming 1 mol of a compound from its elements in their standard states under standard conditions
73
What is the heat of formation for an element in its standard state i.e. O2?
0
74
What is SATP? What is it usually used for?
"standard conditions"
25 degrees C/298K
1 atm
usually for thermodynamics
75
What is STP? What is it usually used for?
"standard temperature and pressure"
0 degrees C/273K
1 atm
usually used for gases
76
Does an exothermic reaction go from strong bonds to weak bonds or vice versa?
an exothermic reaction goes from weak bonds to stronger bonds
77
What is the second law of thermodynamics?
entropy of the universe is always increasing
| the disorder of the universe increases in a spontaneous process
78
How do you calculate entropy? When is it favourable?
Sproducts - Sreactants
| favourable when positive
79
What is the third law of thermodynamics?
defines absolute zero to be a state of zero-entropy
thermal energy is absent, and only the least energetic thermodynamic state is available to the system
0K
80
What is Gibbs free energy?
the energy associated with a reaction that can be used to do work
81
How do you find Gibbs free energy?
delta G = delta H - (TdeltaS)
Note: H and S need to be in the same units and T needs to KELVIN
(when standard needs to be at 298K)
82
When is a reaction exergonic? Endergonic?
Exergonic when delta G 0
83
Is activation energy the same for forward and reverse reactions?
no
| but G, H and S are (opposite sign)
84
What do physical changes depend on?
intermolecular forces
85
What has stronger intermolecular forces a solid or a gas?
solid
86
What is deposition?
gas to solid
87
What is sublimation?
solid to gas
88
What is vaporization?
liquide to gas
89
What is the difference between temperature and thermal energy?
temperature is a measure of average kinetic energy
| thermal energy is a measure of total kinetic energy i.e. depends on the amount
90
When there is a transfer of thermal energy what can happen to a compound?
change of phase OR change in temperature BUT NOT BOTH AT THE SAME TIME
91
How do you find the amount of heat absorbed or released in a phase change?
q=n x deltaH (heat of transition in J/mol or kJ/mol)
+q = heat absorbed
-q = heat released
92
How do you find the amount of heat absorbed or released in a temperature change?
q=mcdeltaT
93
What is the difference between specific heat capacity (c) and heat capacity (C)?
specific heat capacity is for compounds or molecules, use q=mcdeltaT
heat capacity is for objects, use C=mc
94
What is greater, deltaH vaporization or deltaH fusion?
vaporization because to go from liquid to gas you need to break IMFs, but to go from solid to liquid you just need to weaken IMFs
95
What does the slope of a phase transition diagram give?
1/C i.e. you can find C using it (C=1/slope)
96
What do phase diagrams show?
they show the dependence of state of a particular substance on pressure and temperature
97
GIve the relative positions of the different staters in a phase diagram
solid is on the left/top
liquid is on the top middle
super-critical fluids are on the top right
gas is on the bottom
98
What is a super-critical fluid?
has both liquid and gas properties
99
What is the triple point of a phase diagram?
where all 3 phases are in equilibrium
100
What is the critical point of a phase diagram?
the end of the liquid gas boundary (then you start to get super-critical fluids)
101
What is different about H2Os phase diagram compared to most substances?
liquid water is more dense than solid water so at elevated pressure the liquid state is favoured (freezing point is depressed at elevated pressures)
102
What is pressure equal to?
P=force/area
103
What are the assumptions of kinetic molecular theory?
1) volume of gas particles is negligible compared to the volume of the container
2) collisions are elastic (NO IMFs present)
3) particles will travel at a variety of speeds but the average kinetic energy is directly proportional to the temperature
104
How many litres in a cubic metre?
1000L = 1m^3
105
What is the ideal gas constant?
0.0821 Latm/molK
106
Under what conditions do gases behave more ideally?
high temperature and low pressure
107
Under what conditions do gases not behave ideally i.e. follow KMT?
low temperature and high pressure
108
Why doe gases not follow KMT at low temp and high pressure?
when the pressure is high the volume of the particles isn't negligible
when temperature is low there are intermolecular forces
109
Are the pressure and volume (of free space) of real gases greater or less than those of ideal gases?
pressure and volume of free space of real gases are LESS than the pressure and volume of free space in ideal gases
"real world is less than ideal"
110
Are the molecular weights of most ideal gases small or large?
small
111
What is van der Waals equation for?
corrects for the IMFs and volume of particles of a real gas
112
What part of van der Waals equation corrects for IMFs?
an^2/v^2
| the part added to P
113
What part of van der Waals equation corrects for the volume of particles?
nb
| the part subtracted from V
114
What is the combined gas law?
P1V1/T1 = P2V2/T2
115
What is Boyle's law?
P1V1=P2V2 at constant T
| ie pressure is indirectly proportional to volume
116
What is Charles' law?
"CTV"
V1/T1 = V2/T2 at constant P
ie volume is directly proportional to temperature
117
What does T need to be in in order to use any of the gas laws?
Kelvin
118
What is Avogadro's law?
for 2 gases at the same V, T and P the number of moles will also be the same
"standard molar volume"
119
What is standard molar volume at STP?
22.4L (volume of 1 mol of ideal gas at STP)
120
Do heavier gases travel faster or slower than lighter gases?
slower
121
When vapour pressure is equal to atmospheric pressure, what is boiling point equal to?
temperature
122
Is vapour pressure affected by external pressure? What about boiling pressure?
vapor pressure is NOT affected but boiling point IS (when atmP decreases, boiling point decreases)
123
What is effusion?
the escape of gas molecules through a VERY small hole into a vacuum
124
What is Graham's law of effusion?
rate of effusion A/ rate of effusion B = square root (MMB / MMA)
125
What are Arrhenius acids and bases?
acids- produce H3O+ in water
| bases- produce OH- in water
126
What are Bronsted-Lowry acids and bases?
acids- H+ donors
| bases- H+ acceptors
127
What are Lewis acids and bases?
acids- electron acceptors
| bases- electron donors
128
What are the strong acids you need to know for the MCAT?
```
HCl
HBr
HI
HClO4
HNO3
H2SO4
```
129
What are the strong bases you need to know for the MCAT?
Group 1 hydroxides ie NaOH, KOH etfc
Group 1 oxides ie LiO2
Some group 2 hydroxides: Ba(OH)2, Sr(OH)2, Ca(OH)2
Metal amides ie NaNH2
130
With what Ka/Kb is an acid or base considered strong?
1
131
What is Kw at 25 degrees C?
1 x 10^-14
132
How do you find Ka or Kb using Kw?
Ka x Kb = Kw
133
How do you approximate pH on the MCAT?
[H+] = y x 10^n
pH is between n and n-1
large y = closer to n-1
small y = closer to n
134
What does log1 equal?
0
135
What is the Henderson Hasselbalch equation?
pH=pKa+log(base/acid)
136
To maintain a pH of 4 what pKa or pKb of buffer would you want?
pKa=4
| pKb=10
137
What is the Henderson Hasselbalch equation for a base buffer?
pOH=pKb+log(acid/base)
138
What happens when you increase the amount of buffer in a solution?
the buffering capacity increases
| Note: you still need equal amounts of acid and base though
139
What cations make neutral salts?
Na+, K+, Cl-
140
What cations make acidic salts?
```
NH4+
Be2+
Cu2+
Zn2+
Al3+
Cr3+
Fe3+
```
141
What anions make negative salts?
F- etc
142
If you have both acidic and basic ions in a salt solution, how do you know if it acidic or basic?
Ka>Kb, acidic
| Kb>Ka, basic
143
What is the equivalence point on a titration curve?
the point where all acid and base have reacted
144
How do you know what the pH is at the equivalence point?
strong one wins out i.e. if you're adding strong acid then its acidic and vie versa for bases
145
What is the 1/2 EP on a titration curve?
exactly half the weak acid or base has been converted to its conjugate
pH = pKa
146
What is the buffering region on a titration curve?
where you have similar amounts of weak acid or base and conjugate
change in pH is slow here
147
What is different about a titration curve for a strong acid and strong base being titrated?
doesn't have a buffering region
148
What pH do you want an indicator to be?
the pH of the equivalence point
149
What is oxidation?
losing electrons
150
What is reduction?
gaining electrons
151
Is a positive or negative reduction potential spontaneous?
positive
152
What is a galvanic cell?
uses a spontaneous redox reaction to generate an electric current
(also called voltaic cells)
153
Where does oxidation occur?
always at the anode
154
Where does reduction occur?
always at the cathode
155
In which direction do electrons flow?
always from anode to cathode
156
Where do anions and cations migrate?
anions migrate to the anode, cations migrate to the cathode
157
Is the anode in a galvanic cell negative or positive?
negative
158
Is the anode in an electrolytic cell negative or positive?
positive
159
What does F stand for?
faraday
the charge of one mole of electrons
96 500C
160
What is the equation relating delta G to cell voltage?
delta G = -nFEknow
161
If cell voltage is positive is the reaction spontaneous?
yes
162
What is the Nerst equation?
shows how temp and concentration of reactants can alter the voltage of a reaction
E = Eknow - RT/nFln(Q)
163
What is a concentration cell?
galvanic cell that has identical electrodes, but which has half-cells with different ion concentrations so they have a potential difference and a current is produced
164
Where do electrons flow in a concentration cell?
flow to the half-cell with the higher concentration of positive ions
165
What colour is Ce4+? Ce3+?
Ce4+ is bright yellow
| Ce3+ is colourless
166
What is an electrolytic cell?
uses an external voltage source i.e. a battery to create an electric current that forces a non spontaneous redox reaction to occur
167
Are rechargeable batteries galvanic or electrolytic?
galvanic when being discharged
| electrolytic when recharging
168
What are the 2 equations used for electrolysis?
```
Q= It
Q= nF
```
169
When does the colour change occur in a redox titration?
at the equivalence point
170
What is E at the 1/2 EP of a redox titration?
E= standard E of reduction of the species being oxidized (analyte)
171
What is E at the 2x EP of a redox titration?
E= standard E of reduction of the species being reduced
172
What is alpha decay?
Emit an alpha particle which is 2 p+ and 2 n0 (a He nucleus)
| it is the least harmful because alpha particles are relatively large
173
What is beta minus decay?
conversion of a neutron to an e- and p+ (beta- particle is an electron)
atomic number increases, atomic mass stays the same
174
What is beta plus decay?
conversion of a p+ to a e+ and n0
e+ is a positron (beta + particle)
atomic number decreases, atomic mass stays the same
175
What is the 3rd type of beta decay?
e- capture
conversion of an e- and p+ to an n0
atomic number decreases, atomic mass stays the same
176
Are alpha or beta particles more harmful?
beta because they are smaller and have more energy
177
What is gamma decay?
emit energy only, no mass, no charge
most powerful and most harmful
mass and atomic numbers stay the same
178
What is nuclear binding energy?
amount of energy released when protons and neutrons combine to form a nucleus
the more energy released the more stable the nucleus is
179
What happens to mass when a nucleus forms?
mass decreases because some is converted to energy
"mass defect"
Eb=delta m x c^2
180
What is Eb measured in?
MeV, 931.5MeV = 1amu
| to find nuclear binding energy multiply the change in mass in amu by 931.5 MeV
