Gen Chem 1 Flashcards
1
Q
What is the atomic number (Z)
A
the number of protons an atom has
2
Q
What is the mass number (A)
A
the combined number of protons and neutrons
3
Q
When the convention showing the atom, it’s atomic number and mass number how is it arranged
A
the mass number on top of the atomic number. both next to the atom name
4
Q
What are valence electrons
A
electrons farthest away from the nucelus
5
Q
what happens to valence electrons as they get farther away from the nucleus
A
the farther away they are the less they are attracted to the nucleus, and the more they can interact with other atoms.
6
Q
What is an ion
A
a positive or negative atom due to the loss or gain of electrons
7
Q
how many, n’s, p’s, and e-‘s in a Nickel-58 atom
A
Ni = 28 protons
58-28 = 30 neutrons
since it is neutral it has 28 electrons too
8
Q
How many n’s p’s, and e-‘s does a +2 Nickel-60 atom have
A
Ni = 28 protons
60-28 = 32 Neutrons
Since it is plus 2 it has lost 2 electrons = 26 e-‘s
9
Q
How much is one amu in grams
A
1.66 x 10^-24
10
Q
How many amu does carbon have
A
12
11
Q
what is the atomic weight
A
the weight in grams of one mole of a given element (g/mol)
12
Q
What is avogadro’s number
A
the number of particles in a mole
6.022 x 10^23
13
Q
What are isotopes
A
when atoms (having the same number of protons) have different numbers of neutrons
14
Q
how does isotopicity affect reactivity
A
it doesn’t change a lot because they still have the same number of protons and e-‘s
15
Q
How is the atomic weight of the elements on the periodic table found
A
you take the mass number of each of the isotopes multiplied by the frequency of that isotope, and add all of those together
16
Q
Q has
isotope A 60% at 40 amu
isotope B 25% at 44 amu
isotope C 15% at 41 amu
A
40 x .6 = 24
44 x .25 = 11
41 x .15 = 6.15
24 + 11 + 6.15 = 41.15 g/mol
17
Q
What is plancks constant
A
6.626 x 10^-34 J*s
18
Q
what is the equation for the energy of a quantum
A
E = hf h = plancks constant f = frequency
19
Q
What is the equation for the angular momentum of an electron
A
angular momentum = nh/2pie
n = quantum number
h = plancks constant
20
Q
What is bohrs model of the Hydrogen atom
A
a single proton with an electron circling it
21
Q
Since angular momentum of an electron has thee different constants in the equation, what determines angular momentum of an electron
A
the quantum number
22
Q
What is the equation for the energy of an electron
A
E = -Rh/n^2 Rh = 2.18 X 10 ^-18 j/electron
23
Q
What is Rh
A
the rydberg constant = 2.18 x 10^-18 j/electron
24
Q
What is the only factor that affects the energy of an electron
A
the quantum number
25
What will the energy of an electron be in any of it's quantized states
negatve
26
The energy of an electron is related to
it's orbital radius the larger the radius the more energy it has. the smaller the radius the less energy it has
27
what is the smallest orbit an electron can have
n = 1
28
What is the ground state level for an electron
when the electron is at it's lowest energy state
29
How is atomic emission specrta gathered
you take an atom in its ground state, then heat it up or add energy to it and the electron is excited, or moved to a higher energy state. It will quickly move back to it's ground state but in doing so it will release energy in the form of photons.
30
What is the equation for the energy of the photons released when doing atomic emission spectra
```
E = hc/Lambda
H = plancks constant
c = speed of light
Lambda = wavelength
```
31
What is the speed of light
3 x 10^8 m/s
32
How are atomic emission spectra used to identify an element
each electron in a element when returning to ground state releases a specific wavelength of light, each element has its own specific electrons so each element has it's own unique atomic emission spectrum
33
Can the atomic emission spectrum of stars be done
ye[p
34
What is the balmer series of hydrogen emission lines
transitions from n > 2 to n = 2
35
how many wavelengths are in a balmer series of hydrogen emission lines
4 in the visible region
36
what is the lyman series of hydrogen emission lines
transitions fron n > 1 to n = 1
37
what causes an electron to jump up to a higher energy state
it absorbs energy
38
what is an absorption spectrum
it is the spectrum showing which wavelengths of energy are absorbed by a specific element
39
When can absorption spectrum be used
in the identification of elements present in a gas phase sample
40
Why does Bohr's model not work for atoms with more than one e-
because it doesn't take the repulsion between each electron into account
41
What is the main difference between Bohr's model and modern quantum mechanics
that electrons don't go around the nucleus in a circle, but they move around in orbitals
42
what does an electron orbital show
it represents the probability of finding an electron in a given region
43
what is the Heisenberg uncertainty principle
states that it is impossible to determine with perfect accuracy the momentum and position of an electron simultaneously
44
What are the four quantum numbers
n, ℓ, mℓ, ms
45
What is the pauli exclusion principle
no two electrons in a given atom can posses the same set of four quantum numbers.
46
What does the position and energy of an electron described by its quantum numbers give you
its energy state
47
What is the quantum number n and what does it indicate
the principal quantum number
| it indicates the distance of the electrons from the nucleus
48
What is the quantum number ℓ and what does it indicate
the azimuthal quantum number (angular momentum quantum number)
it refers to the subshells of the valence electrons
49
```
what do the ℓ quantum numbers
0
1
2
3
indicate
```
0 means that the electrons are in the S subshell
1 means the p subshell
2 = d subshell
3 = f subshell
50
With using ℓ, how many electrons can fill particular subshells
4f + 2
51
the greater the number of ℓ, the
greater the energy of the subshell
52
what is the quantum number mℓ, and what does it indicate
it is the magnetic quantum number
| it indicates which orbital
53
how many mℓ values can there be in terms of ℓ
2ℓ + 1
54
What is the ms quantum number and what does it indicate
it is the spin quantum number
| and it is just the last way to identify electrons (both electrons in one orbital must have opposite spins
55
What are electrons with parallel spins
electrons in different orbitals with the same Ms values
56
What are paired electrons
electrons in the same orbital with opposite spins
57
In the electron configuration what does the
1st number
letter
2nd number indicate
1. the principal energy level
2. the subshell
3. the number of electrons in that subshell
58
how do you know which subshells will fill first
those with a lower n + ℓ value, if they are tied, then you go with the one with the lower n value
59
In what order are subshells filled
```
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
6s
```
60
What is hund's rule
all orbitals of the same energy level fill halfway before any will completely fill
61
What are some exceptions to the orbital filling rules
Chromium
- expected = 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d4
- actual = 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5
This is so that the 3d orbitals can be half filled
Copper
- does the same thing except 5 electrons later, one e- is promoted to the 3d orbital from the 4s obital to give it 10 e-'s
62
what are paramagnetic materials
materials with unpaired electrons, that are weakly attracted to a magnetic field
63
What are diamagnetic materials
materials with no unpaired electrons that are slightly repelled by a magnetic field
64
What are valence electrons
all of the electrons in it's outer energy shell
65
What is the periodic law
that the chemical properties of the elements are largely dependent on their atomic number
66
What are the periods of the p table
rows
67
what are the groups on the p table
columns
68
what does the roman numerals on the periodic table groups indicate
the number of valence electrons
69
What does the A or B indicate after the roman numeral
```
A = representative elements (either s or p as outermost orbitals)
B = nonrepresentative elements ( have d as their outermost orbitals)
```
70
What are the two kinds of non-representative elements
transition elements
| lanthanide and actinide serie
71
what are transition elements
those with partially filled d orbitals
72
what are lanthanide and actinide elements
those with partially filled f orbitals
73
what do all elements strive for
8 valence electrons
74
how does atom size radii across the p table
as you go from left to right they get smaller
| as you go from bottom to top they get smaller
75
how do they find atomic radii
it is one half of the distance between to atoms of that element that are just touching each other
76
why does atomic radii get smaller towards the top right of the p table
from left to right because the more protons and electrons you have the more that they are pulled together
from bottom to top because filled orbitals block the attraction for the valence electrons and the nucleus.
77
what is the ionization energy
the amount of energy required to remove an electron from an atom
78
how does ionization energy change on the p table
the same way that atomic radii does. it is the largest in the top right because the electrons are the most tightly held together
79
what is first and second ionization energy and how do they compar
first ionization energy describes the energy required to remove the first electron, second " " second electron
second ionization energy is always greater than the first ionization energy
80
What is electron affinity
the energy change that occurs when an electron is added to a gaseous atom, (it represents the ease with which an atom can accept an electron)
81
how does atomic size affect electronegativity
the smaller the atom, the more it will pull on the electron
82
how does electronegativity change across the p table
it increases towards the top left
83
what is the effective nuclear charge (Zeff)
the attractive pull of the nucleus (increases from left to right)(more protons)
84
What is electronegativity
the measure of attraction an atom has for electrons in a chemical bond.
85
what is special about cesium
-largest
- most metallic
least electronegative of all naturally occurring elements
- smallest ionization energy
- least exothermic electron affinity
86
What is another way to think about electronegativity
nuclear positivity
| the stronger the nuclear pull the more electronegative it will be
87
What is the most electronegative atom
fluorine
88
What are the three types of elements
metals
non-metals
metalloids (semimetals)
89
where are metals, non-metals, and mettaloids located on the p tabel
metals left of the ziggy line
non-metals on the right of the ziggy line
mettaloids form the ziggy line
90
What are the characteristics of metals
- shiny solids at room temperature (except mercury)
- high melting points
- high densitties
- malleable
- ductile
- low ionization energies
- low electronegativity
91
What is malleability
the ability of a metal to be hammered into shapes
92
What is ductility
the ability to be drawn into wire
93
What gives rise to the characteristic properties of metals
the few electrons in the valence shells can easily be removed
94
What makes metals good conductors of heat and electricity
that their electrons can easily be moved.
95
What are some of the characteristics of non-metals
- brittle in solid state
- little or no luster
- high ionization energies and electronegativities
- gain electrons easily
96
What are some of the characteristics of metalloids
- posses characteristics of both metals and non-metals
97
What elements are metalloids
```
Boron
Silicon
Germanium
Arsenic
Antimony
Tellurium
(polonium)
```
98
What are alkali metals
elements in group IA
99
What are the characteristics of alkali metals
- one loosely bound e-
- highly reactive due to low ionization energies
- react highly with halogens
100
What are alkaline earths
elements in group IIA
101
What are some of the characteristics of alkaline earths
- large atomic radii
- fairly loosely bound electrons 2
- low electronegativity, ionization energies, and electron affinities
102
What are halogens
elements in group VIIA
103
What are the characteristics of halogens
- 7 valence electrons
| - high electronegativities
104
What are the physical states of halogens at room temperature
F and Cl are gasses
Br is a liquid
I is a solid
105
What are the noble gases
the elements of group VIIIA
106
what are the charateristics of the noble gasses
- quite unreactive
- high ionization energies
- low boiling points
- all gasses at room temp
107
What are the transition elements
elements in group IB to VIIIB
108
What are the charateristics of transition elements
- very hard
- high melting points
- malleable
- highly electrically conductive
- low ionization energies
109
What are the positively charged transition metals called
oxidation states
110
What causes transition metals to be malleable and conductive
they have loosely held d-orbital electrons
111
What is it called when transition metals in their oxidation states complex with water molecules, or nonmetals
```
hydration complex (with water)
complexes without
```
112
What happens with complexes
they absorb a wide variety of frequencies of lights. which leads to interesting colors of liquids.
113
What interacts to form bonds between molecules
the valence electrons of molecules
114
what is the difference between intramolecular forces and intermolecular forces
```
intra = the forces that actually hold the two atoms together
inter = other forces within the molecule that are between atoms that aren't bonded
```
115
What is the octet rule
that atoms bond to other atoms until they have 8 electrons in their outermost electron shell
116
what are exceptions to the octet rule
```
Hydrogen (2 or 4)
Lithium and beryllium (2 or 4)
Boron (6)
elements beyond the second row
phosphorus and sulfur (8+)
```
117
What is ionic bonding
when electrons from an atom with a small ionization energy are transferred to an atom with a high electron affinity.
118
what kinds of forces hold ionic bonds together
electrostatic
119
what is covalent bonding
When an electron pair is shared between two atoms.
120
What is it called when the bonds are partially covalent and partially ionic
polar covalent bonds
121
in which kind of bond are electrons completely donated, what kind are they just shared
donated in ionic
| shared in covalent
122
how large does the difference in electronegativity need to be for it to be a strong enough ionic bond to create cations and anions
at least 1.7 difference or larger
123
what are the characteristics of ionic compounds
- high boiling and melting points
- conduct electricity in the liquid and aqueous states (not in solid states)
- form crystal lattices that increase attractive forces and decrease repulsive forces
124
What kind of bond is formed when elements of similar electronegativities bind
covalent bond
125
which is requires more energy. the creation of two ions, or the energy released when an ionic bond is broken
the energy required to create ionic bonds
126
what are the characteristics of covalent compounds
- weak intermolecular forces
- low melting solids
- do not conduct electricity in liquid or aqueous states
127
what is bond order
it tells you how many electron pairs are shared in a bond
single bond = bond order of 1
double bond = bond order of 2
...
128
What are the two features of a covalent bond
bond length
| bond order
129
how are bond order and bond length related
higher bond orders = shorter bond lengths
130
What is bond energy
the amount of energy required to seperate two bonded items
131
how is bond order and bond energy related
the higher the order the higher the bond energy
132
what is a formal charge
when more or less electrons are attributed to an atom in a molecule than would be alone
133
What are the steps of drawing lewis structures
1. Draw the skeletal structure of the molecule
2. count all of the valence e-'s (sum = total number of electrons)
3. draw single bonds
4. complete all of the octets (outside atoms first)
5. put extra electrons on the central atom
6. double or triple bonds if needed
134
How do you calculate the formal charge of an atom
1. take the number of valence electrons it normally has
2. subtract half of all the bonding electrons
3. subtract all of the non-bonding electrons
135
What is a resonance structure
when two or more non-identical lewis structures can be drawn
136
how do we know which resonance structure is the actual molecule
it's not one or the other but a hybrid of both, it spends more time, or is more like the more stable of the two resonance structures
137
What kind of arrow represents resonance structures
one double headed arrow
138
Are lewis structures with formal charges preferred over lewis structures without formal charges
no structures without formal charges are more preferred
139
What are some exceptions to the octet rule
atoms found in or below the third period can have more than 8 valence electrons because the D orbital can hold some
140
What are polar covalent bonds
bonds formed between atoms with small electronegativity (.4-1.7)
the electron pair is held more closely to the more electronegative atom
and thus it has a partial negative charge and the other has a partial positive charge
141
what is a dipole moment
the product of charge magnitude and the distance between the partial charges in a polar covalent bond
142
What are coordinate covalent bonds
when the bonding electron pair comes from only one of the two atoms
143
where are coordinate covalent bonds normally found
in lewis acid-base complexes
144
what is a lewis acid
| what is a lewis base
lewis base = electron pair donor
| lewis acid = electron pair acceptor
145
2 regions of electron density
| shape and degree
linear
| 180 degrees
146
3 regions of electron density
trigonal planar
| 120 degrees
147
4 regions of electron density
tetrahedral 109.5 degrees
148
5 regions of electron density
trigonal bipyrimidal
| 90, 120, 180 degrees
149
6 regions of electron density
octahedral
| 90, 180 degrees
150
What does an ℓ=0 orbital look like
it is an s-orbital
| a perfect sphere
151
What does a ℓ=1 orbital look like
that is a p-orbital
| there are three along the x,y,and z axis. they are two baloons tied together shape (dumbell)
152
What does a ℓ=2 orbital look like
That is a d orbital
| you don't need to know
153
What is a molecular orbital
when two atomic orbitals interact during a bond
154
What is the molecular orbital of a single bond
both of the orbitals overlap head to head and form sigma bond
155
What is the molecular orbital of a double bond
the orbitals are parallel and it is called a pie bond
156
How do you get a bonding orbital
when the two bonding atomic orbitals are the same sign, they bond
157
how do you get a non-bonding orbital
when the two bonding orbitals are opposite signs, they don't bond
158
What are intermolecular forces
forces between molecule
159
What are the kinds of intermolecular forces
Dipole-dipole interactions
hydrogen bonding
dispersion forces
160
What are dipole dipole interactions
when the polar (+) end of one molecule orients itself near the polar ( - ) end of another molecule
161
when do dipole dipole interactions occur
in solid and liquid phases
| gasses the molecules are too spread out
162
what can dipole-dipole interactions do for a solid or molecule
they can increase the melting and boiling points compared to non-polar substances of the same weight
163
What is hydrogen bonding
when a H bonded to an electronegative atom (F,O,N) and it's partial positive charge interacts with the partial negative charge on other molecules
164
what does hydrogen bonding do to the melting point and boiling points of molecules
raises them significantly
165
what kinds of molecules have greater dispersion forces
large atomed molecules
166
what determines if molecules are miscible or immiscible
the intermolecular forces
167
what is a compound
a pure substance that is composed of two or more elements in a fixed portion
168
What is the ionic substitute for molecular weight
formula weight
169
What is the Gram Equivalent weight
the molar mass (how many grams per mole a substance is)
| divided by the number of (hydrogens) that the substance can supply.
170
how do you find equivalents
weight of the compound / Gram equivalent weight
171
what good can equivalents of weight do
it can help you know how much acid you will need to neutralize a base
172
What is the law of constant composition
states that any sample of a given compound will contain the same elements in the same ratio
173
What is the empirical formula
the simplest whole number ratio of the elements in a compound
174
what is the molecular formula
gives the exact number of atoms of each element in a compound
175
What is the empirical then molecular formula of benzene
```
empirical = CH
molecular = C6H6
```
176
What is the percent composition
the weight percent of a given element in a compound
177
How do you find percent composition
get the mass on the substance in the formula / the formula weight of the compound
178
what are combination reactions
two or more reactants form one product
179
what are decomposition reactions
one reactant breaks down into two or more products
180
What are single displacement reactions
Redox reactions
| when one atom or ion moves from one compound to the other
181
What is a double displacement reaction
a metathesis reaction
| when elements of two compounds react to form two different compounds
182
what is a spectator ion
an ion in an ionic equation that doesn't change from reactants to products
183
what is a net ionic equation
an ionic equation in which the spectator ions are left out
184
What are neutralization reactions
a double displacement reaction in which water and a salt are created and an acid and base are neutralized
185
What is the mechanism of a reaction
the actual steps through which the reaction occurs
186
What is the rate determining step
the slowest step of the reactoin
187
what are the two ways to determine the rate of a reactoin
the dissapearance of reactants/time
| the formation of products/time
188
how do you distinguish between rates by loss of reactants vs. formation of products
rate by loss of reactants will be negative
189
what are the units of reaction rates
moles/L*s
| M/s
190
how do you find the exponents in the rate law equation
they are equal to the coefficients of the substances IN THE RATE DETERMINING STEP
191
how do you find the reaction order for a reaction
add the exponents of the rate law
192
What is a zero order reaction
one in which the rate is independent of the concentrations of the reactans
193
what is a first order reaction
the rate is proportional to the concentration of one reactant
194
what is the radioactive decay equation
```
At = A0 e^-kt
At = concentration at time t
A0 = initial concentration
k = rate constant
t = time
```
195
what is the half life equation
t1/2 = .693/k
196
What is a second order reaction
the rate is proportional to the product of the concentration of two reactants, or to the square of one reactant.
197
What are the units for a zero oder reaction
Ms-1
198
What are the units for a first order reaction
s-1
199
What are the units for a second order reaction
M-1s-1
200
What are higher order reactions
reactions that have a reaction order greater than 2
201
what are mixed order reactions
reactions with a fractional order (X^1/3)
202
What is the collision theory of chemical mechanics
- the rate of a reaction is proportional to the number of collisions per second between reacting molecules
203
What are effective collisions
collisions that result in formation of a product
204
what has to happen for a collision to be effective
the molecles must be moving fast enough and be properly oriented
205
What is the activation energy
the minimum amount of energy needed to break old bonds and form new ones
206
how do you get the rate of reactions with total number of collisions per second (Z) and the fraction of collisions that are effective (f)
rate = fZ
207
What is a transition state
when the old bonds are beginning to break and the new bonds are beginning to form. the transition state then degrades into the product
208
What is the energy level of the transition state like
it is higher than either the products or the reactants (those use energy from a collision to get to that energy)
209
what is the difference between an intermediate and a transition state
transition states have a very very short finite lifetime
210
what does ΔH refer to
enthalpy change
| this shows the difference between the potential energy of the reactants and the potential energy of the products
211
What does a - or + ΔH mean
-- means that it is exothermic (heat is given off)
| + means that it is endothermic (heat is absorbed)
212
What two things can increase the reaction rate
more effective collisions
| stabilization of the activated complex (transition state)
213
What four things affect reaction rates
1. concentration
2. temperature
3. medium
4. catalysts
214
When have we reached equilibrium
when there is no net change of the concentrations of products and reactants during a reversable reaction
215
What is the law of mass action
at equilibrium the forward and reverse reaction rates will be equal
[C]c[D]d / [A]a[B]b = Kc
216
What is Kc
the equilibrium constant (except in gasses it is Kp)
217
What is the reactant quotient Q indicate
Q is a measure of the degree to which a reaction has gone to completion
218
what is Q equal to
[C]c[D]d / [A]a[B]b
219
What are the characteristics of the equilibrium constant Keq
- pure solids and liquids don't count
- if Keq is very large when compared to 1 it is mostly products left
- if Keq is very small compared to 1 is it mostly reactants left
- if Keq is close to 1 they they are at equilibium, and the concentrations of products and reactants will be very similar
220
What happens to reactions as you increase the pressure
the reaction will go to whatever favors less moles, and decrease volume
221
What is an isolated system
it cannot exchange energy or matter with the surroundings
222
what is a closed system
energy, but not matter can be exchanged with the surroundings
223
what is an open system
energy and matter can be exchanged with the surroundings
224
what is an isothermal process
the process occurs while the temperature remains constant
225
What is an adiabatic process
when no heat exchange occurs
226
what is an isobaric process
then the process occurs while the pressure remains constant
227
What is negative and positive heat
positive heat is heat absorbed by the system
| negative heat is heat lost by the system
228
What are endothermic reactions
those that absorb heat
229
what are exothermic reactions
those that lose heat
230
how many joules are equal to one calorie
4.184 J
231
what is the equation for finding the heat (q) lost or gained by a system
```
q = mcΔT
q = heat
m = mass
c = specific heat
ΔT = temperature change
```
232
in constant volume calorimetry how do you find the q or the reaction
q rxn = - (q water + q steel)
233
What does the little circle after enthalpy, entropy, and free energy mean
that it is done at standard conditions (25 degrees celcius, 1 atm)
234
What is enthalpy (H)
it accounts for the heat absorbed or lost by the system
235
How do you find the enthalpy change
ΔH rxn = H products - H reactants
236
what does a +ΔH represent
a endothermic reaction (absorbs energy)
237
What is Hess's law
the enthalpies of reactions are additive
238
What is the bond dissociation energy
the amount of energy needed to break a particular type of bond in one mole of gaseous molecules
239
how do you use bond energy to calculate ΔHrxn
ΔHrxn = (ΔH bonds broken) + (ΔH bonds formed)
240
What is the bond dissociation energy
the amount of energy needed to break a particular type of bond in one mole of gaseous molecules
241
how do you use bond energy to calculate ΔHrxn
ΔHrxn = (ΔH bonds broken) + (ΔH bonds formed)
242
What are some rules of entropy
- S solid < S liquid < S gas
| - all spontaneous processes increase the entropy of the system
243
how do you find ΔS
```
ΔS = qrev/T
ΔS = Sfinal-Sinitial
```
244
When has entropy reached it maximum state
at equilibrium
245
how do you find ΔS
```
ΔS = qrev/T
ΔS = Sfinal-Sinitial
```
246
what does Gibbs Free Energy tell you
the spontaneity of a reaction
247
What is the equation for Gibbs Free energy
ΔG = ΔH - TΔS
248
```
What happens if ΔH and ΔS are...
+ +
+ -
- +
- -
```
+ + = spontaneous at high temperature
+ - = non spontaneous
- + = spontaneous at all temperatures
- - = spontaneous at low temperatures
249
What does ΔG = when considering Keq
```
ΔG = -RT ln Keq
ΔG = -RT ln Q
ΔG = ΔG + RT ln Q
```
250
What does ΔG = when considering Keq
```
ΔG = -RT ln Keq
ΔG = -RT ln Q
ΔG = ΔG + RT ln Q
```
251
how many mmHG is 1 atm
760
252
how many torr is 1 atm
760
253
What are standard conditions
25 degrees celcius
| 198 K
254
What is the equation for Boyle's law (ideal gas)
P1V1 = P2V2
| constant temperature
255
What is the equation for charles law (ideal gas)
V1/T1 = V2/T2
| constant pressure
256
What is avogadros law (ideal gas)
n1/V1 = n2/V2
| constant temp and pressure
257
What is the ideal gas law
PV = nRT
258
What are the two values for R in the ideal gas law
.0821 L*atm/(mol*K)
| 8.314 J/(K*mol)
259
How do you know which R to use in the ideal gas law
look at the units
.0821 (L*atm)
8.314 (J)
260
How do you find density with the ideal gas law
n = m/MM
mass in grams over molar mass
so density = P(MM)/RT
261
What is the volume of one mole of gas at standard temperature and pressure
22.4 L
262
how are PV and T related
```
P1V1/T1 = P2V2/T2
1= STP
2 = actual conditions
```
263
how do you find density of a gas with the ideal gas law
```
1. find V2
V2 = 22.4 (P1/P2)(T2/T1)
2. d =m/V2
divide mass (g) by V2
```
264
how do you find molar mass of a gas with the ideal gas law
```
1. find Vstp
Vstp = V (P2/P1)(T1/T2)
2. find density
g/L (Vstp)
3. Molar mass = d*V1 (V1 = 22.4)
```
265
What are causes of deviation to the ideal gas law
1. High pressure
| 2. Low temperature
266
What is daltons law of partial pressures
the total pressure of a gas is equal to the sum of the partial pressures in the gas
267
how do you calculate the partial pressures of gasses with moles and pressure
```
Pp = (P)(xa)
xa = number of moles of said gas/total moles of gas
```
268
What are the assumptions of the kinetic molecular theory of gasses
1. the size of particles in gasses are negligible compared to volume
2. atoms and molecules don't exhibit intermolecular attractions
3. particles are in continuous random motion, and collisions
4. collisions are elastic
5. average KE is proportional to the average temperature
269
What does KE for an ideal gas =
KE = 3/2 kT
270
how are rates of diffusion related to molecular mass compared
r1/r2 = square root of MM2/MM1
271
how are the rates of effusion related to molecular mass compared
the same way as diffusion
272
what are amorphous solids
has no ordered arrangement of molecules
273
What are the two kinds of crystals
metal and ionic
274
what are the characteristics of ionic crystals
- high boiling/melting points
| - low conductivity
275
What are the three types of unit cells
simple cubic
body-centered cubic
face centered cubic
276
what happens to the liquid as some of it evaporates
the liquid cools due to a loss of kinetic energy
277
what is the definition of boiling point
the temperature at which the vapor pressure of the liquid equals the external pressure
278
what is the gibbs function for phase changes
```
ΔG = G (gas) - G (solid) (if going from a gas to a solid)
os G(gas) = G(solid) at equilibrium
```
279
Where is the gas phase usually found in a phase diagram
high temperatures, low pressures
280
What is the liquid phase usually found in a phase diagram
high temperatures, high pressure
281
Where is the solid phase usually found in a phase diagram
low temperatures, high pressure
282
what is the triple point on a phase diagram
the point (pressure and temperature) where all three phases are in equilibrium
283
What is the critical point
the point (temperature and pressure) at which no distinction between liquid and gas can be made
284
what does adding a solute do to the freezing point of water
causes freezing point depression ( freezing point goes down)
285
What is the equation for freezing point depression
```
ΔT(f) = K(f)m
m = molality
K(freezing) = freezing point depression
```
286
What is molality
moles/kg
287
what happens to boiling point as you add solute to water
causes boiling point elevatoin
288
What is the equation for boiling point elevation
ΔT(b) = K(b)m
289
What is the equation for osmotic pressure
```
II = MRT
II = osmotic pressure
M = molarity
R = ideal gas constant
T = temperature (K)
```
290
What is raoult's law
P(a) = x(a)P(i)(a)
The vapor pressure of substance A is equal to the mole fraction of A multiplied by the vapor pressure of A (pure substance)
291
What is solvation or dissolution
dissolving something in solution
292
When does solvation occur
when attractive forces between the solute and solvent are greater than those between the solute particles
293
what is the general rule of solvation
like dissolves like
294
What is the solubility of a substance
the maximum amount of that substance that can be dissolved in a particular solvent and a particular temperature
295
What happens when the maximum amount of a substance has been dissolved in a solvent
the solution is saturated
296
What things are soluble in water
look at page 280 of kaplan
297
What do the roman numerals after metal ions mead
they show how much of a positive charge the metal has
298
What is ferr -ous and ferr -ic
ferrOUS means it is the lesser of two cations of iron
| ferrIC means it is the greater of two cations of iron
299
what is hydride, fluoride, oxide
-IDE means that they are monatomic anions
300
what are oxyanions
anions that contain oxygen
301
what is the difference between nitrite and nitrate
nitr -ITE is the lesser of two anions
| nitr -ATE is the greater of two anions
302
What do hypo and per mean when speaking of anions
Hypo- chlor -ite means it is the least of four anions
chrorite is the 3rd of four anions
chlorate is the 2nd of four anions
per- chlor -ate is the greatest of four anions
303
what can be added in the name in place of hydrogen when one hydrogen has been added to an anion
bi-
like bicarbonate HCO3-
or bisulfate HSO4-
304
when is a solute a strong electrolyte
if is completely dissolves into its constituent ions
305
what is a weak electrolyte
a solute that doesn't completely dissolve into its constituent ions
306
what are non-electrolytes
solutes that don't ionize at all in solution
307
what is percent composition by mass
mass of the solute divided by mass of the solvent
| multiplied by 100
308
what is the mole fraction
the moles of a compound divided by the rest of the moles in the solution
309
What is molarity
moles of solute per liter of solution (SOLUTION)
310
What is molality
moles of solute per kilogram of solvent (SOLVENT)
311
what is normality
the number of gram equivalent weights of solute per liter of solution (SOLUTION)
312
what is the equation for dilution
M1V1 = M2V2
Molarity
Volume
313
what is the ion product (I.P.) also referred to as the Qsp
IP = [An+]m[Bm-]n
314
What is the solibility product constant
Ksp = [An+]m[Bm-]n (in a saturated solution/equilibrium solution)
315
What is the difference in Qsp and Ksp
Qsp refers to the starting ion product
| Ksp refers to the equilibrium ion product
316
If Qsp > Ksp
precipitation will occur
317
if Qsp = Ksp
nothing will change, the solution is at equilibrium
318
if Qsp < Ksp
the solid will dissolve until saturation
319
what does the common ion effect state
that a salt will not dissolve as well in a solution that already has a common ion
320
what are indicators
things that change colors in acids and bases
321
what is litmus paper
an acid/base indicator
322
what is the arrhenius definition of an acid and a base
```
acid = something that produces an H
base = something that produces an OH
```
323
What is the Bronstead lowry definition of an acid and a base
```
acid = proton donor
base = proton acceptor
```
324
what is the lewis definition of an acid and a base
```
acid = electron acceptor
base = electron donor
```
325
What is F- called
fluoride
326
what is HF called
hydrofluoric acid
327
what is Br- called
bromide
328
what is HBr called
hydrobromic acid
329
What is ClO- called
hypochlorite
330
What is HClO called
hypochlorous acid
331
What is ClO2- called
chlorite
332
what is HClO2 called
chlorous acid
333
what is ClO3- called
chlorate
334
what is HClO3 called
chloric acid
335
what is ClO4- called
perchlorate
336
what is HClO4 called
perchloric acid
337
what is No2- called
nitrite
338
what is HNO2 called
nitrous acid
339
what is NO3- called
nitrate
340
what is HNO3 called
nitric acid
341
what is pH
the measure of Hydrogen ion concentration
342
what does pH =
```
pH = -log [H+]
pH = log 1/[H+]
```
343
what is pOH
the measure of hydroxide ion concentration
344
what does pOH =
```
pOH = -log[OH-]
pOh = log 1/[OH-]
```
345
What is Kw
the water dissociation constant
346
what does Kw =
Kw = [H+][OH-} = 10^-14
347
pH + pOH =
pH + pOH = 14
348
how do yo approximate Ka = 1.8 * 10^-5
Ka = 5 -log 1.8
if the -log number is less that 3.2 the number will be closer to 5 than 4 , if it is more that 3.2 the number will be closer to 4 than 5
349
what does the log 10^x =
x
350
what is Ka and what does it indicate
it is the acid dissociation constant, and it measures how much the acid dissociates
351
What is the Kb and what does it indicate
it is the base dissociation constant, and it measures how much the base dissociates
352
What does Ka =
Ka = [H3O+][A-]/[HA]
353
what kind of Ka does a weak acid have
a small one, the weaker the acid, the smaller the Ka
354
What does Kb =
Kb = [B+][OH-]/[BOH]
355
if Ka is large then Kb is
small
356
what is a neutralization reaction
when an acid and a base neutralize each other, forming a salt
357
what is a hydrolysis reaction
when the salt ions react with water to form acid and base
358
What happens when you add strong acid and strong base
you get a salt and water (pH = 7)
359
what happens when you add strong acid and weak base
you get a salt, but no water. and the cation of the salt can react to reform the weak base, but decreases the ph by using up the OH-
360
what happens when you add weak acid and strong base
you get a salt, the anion of the salt reacts to form the weak acid, the pH decreases because the strong base uses up the H+
361
What happens when you add a weak acid and a weak base
if the Kb is stronger than the Ka than the solution is basic
362
What does Ka =
Ka = [H3O+][A-]/[HA]
363
what kind of Ka does a weak acid have
a small one, the weaker the acid, the smaller the Ka
364
What does Kb =
Kb = [B+][OH-]/[BOH]
365
if Ka is large then Kb is
small
366
what is a neutralization reaction
when an acid and a base neutralize each other, forming a salt
367
what is a hydrolysis reaction
when the salt ions react with water to form acid and base
368
What happens when you add strong acid and strong base
you get a salt and water (pH = 7)
369
what happens when you add strong acid and weak base
you get a salt, but no water. and the cation of the salt can react to reform the weak base, but decreases the ph by using up the OH-
370
what happens when you add weak acid and strong base
you get a salt, the anion of the salt reacts to form the weak acid, the pH decreases because the strong base uses up the H+
371
What happens when you add a weak acid and a weak base
if the Kb is stronger than the Ka than the solution is basic
372
what is polyvalence
when an acid or base has more than 1 equivalent of H+ or OH-
373
What is normality the quantity of acidic or basic capacity
the amount of mole equivalents an acid or base will give off
374
What happens in the Henderson hasselbach equation when conjugate acid concentration = weak base concentration
pOH = pKb
| we are halfway to the equivalence point
375
what are amphoteric or amphiprotic species
species that can act as either an acid or a base
376
what are the steps in balancing a redox reaction
1. separate the two half reactions
2. balance all atoms in both reactions (use H2O to balance O, then use H+ to balance the H atoms)
3. add electrons to balance the charges
4. multiply both sides by a number so they both have the same number of electrons
5. add the half reactions, and cancel out items
6. check to see that they are all balanced
377
what is the equivalence point
when the acid equivalents = base equivalents
378
what is the equivalence point for strong acid and strong base
pH = 7
379
how do you calculate the volume added to reach endpoint in a titration
VN (acid) = VN (base)
V = volume
N = Normality
380
what is a buffer
a mixture of weak acid and it's salt or a weak base and it's salt.
381
what do buffers do
they resist changes in pH when acids or bases are added
382
What is the henderson hasselbach equation for a weak acid buffer solution
pH = pKa + log (conjugate base/weak acid)
383
What is the henderson hasselbach equation for a weak base buffer solution
pOH = pKb + log (conjugate acid/weak base)
384
What happens in the Henderson hasselbach equation when conjugate acid concentration = weak base concentration
pOH = pKb
| we are halfway to the equivalence point
385
What are the rules for assigning oxidation numbers
1. oxidation of free elements is zero
2. oxidation of monatomic ion is = to the charge
3. oxidation number of group IA = 1, IIA = 2
4. oxidation number of grooup VIIA = -1 (except when paired with a more electronegative atom)
5. oxidation of H = -1 in compounds with less electronegative elements than H
6. Oxidation number of O = -2, unless paired to a more electronegative atom O = +1, or unless two are paired O = -1
7. sum of oxidation numbers in a neutral compound = 0
386
what are the steps in balancing a redox reaction
1. separate the two half reactions
2. balance all atoms in both reactions (use H2O to balance O, then use H+ to balance the H atoms)
3. add electrons to balance the charges
4. multiply both sides by a number so they both have the same number of electrons
5. add the half reactions, and cancel out items
6. check to see that they are all balanced
387
what types of reactions occur in galvanic cells
spontaneous redox reactions
388
What is the set up of a galvanic cell
1. Two half cells
- One with a negative anode, the other with a positive cathode
2. the anode and cathode are connected to allow for electron flow
3. The anode is in an aqueous common ion solution
4. The cathode is also in a common ion solution
5. there is a salt bridge with uncommon ions to both solutions
389
how does a galvanic cell work.
1. the - anode is oxidized, losing it's positive ion, and sending electrons through the wire
2. those electrons move across the wire to the positive cathode
3. Those electrons on the cathode cause the positive ions in solution to bind to the cathode and the electrons in reduction.
4. ions in the salt bridge come out to counteract the ions added to the anode solution, and replace the ions lost from the cathode solution
390
What are the rules in drawing a cell diagram of a galvanic cell
1. the reactants and products are always listed from left to right in this form
anode / anode solution // cathode solution / cathode
2. single verticle lines represent a phase boundry
3. double lines represent a salt bridge
391
What is the reaction like in an electrolytic cell
non-spontaneous redox reaction
392
What is the reduction potential of a species
the tendency of a species to acquire electrons and be reduced
393
The more positive the reduction potential of a species...
the more likely it is to be reduced
394
what is the standard electromotive force (emf)
the difference in potential energy between two half cells
395
how do you get emf
emf = reduction potential red. + reduction potential cat
396
What is the sign for reduction potential
E (knot)
397
what is the equation for determining spontaneity from emf
ΔG = -nFE(cell)
n = number of moles of electrons exchanged
F - faradays constant
E = emf of the cell
398
What to remember about ΔG and emf
if faradays constant is in j/V, then ΔG must be in j not kj
399
what is the nernst equation
Ecell = E(knot)cell - (RT/nF)(lnQ)
400
what is another way to get ΔG(Knot)
```
ΔG(knot) = -RT ln Keq
R = 8.314 J/ (k*mol)
```
401
nFE(knot)cell = RT ln Keq
nFE (knot)cell = RT ln Keq
| (if nFE is positive then Keq is positive and the reaction moves to the right
402
what is the mass defect
the fact that every nucleus has a smaller mass than the combined masses of protons and neutrons
403
what is the cause of mass defect
that some mass has been converted into energy (binding energy)
404
what atom has the most binding energy
iron
405
how do you find binding energy
find the mass defect and plug it into E = mc^2
406
What is fusion
the combining of small nuclei to form a larger nucleus
407
what is fission
the splitting of a large nucleus to from two smaller nuclei
408
can cause fission
the addition of a neutron
409
What is alpha decay
loss of a 4HE (2 protons, 2 neutrons)
410
what happens in alpha decay
- 4 mass number
| - 2 atomic number
411
What is beta decay
loss of an electron
412
which is more penetrating alpha or beta decay
beta
413
What is alpha decay
loss of a 4HE (2 protons, 2 neutrons)
414
what happens in alpha decay
- 4 mass number
| - 2 atomic number
415
What is beta decay
loss of an electron
416
which is more penetrating alpha or beta decay
beta
417
What are the two type of beta decay
B- (an electron)
| B+ (a positron)
418
What does a negative beta decay do
+1 in atomic number
419
What does positron emission do
-1 in atomic number
420
What is gamma decay
emission of high energy photons
421
what happens in gamma decay
nothing to mass and atomic number, just a loss of energy
422
what is electron capture
gaining an electron that combines with a neutron to form a proton
423
what happens in electron capture
no change to mass number
| +1 atomic number
424
what is the equation for exponential decay
```
n = n(knot) e^-(gamma t)
n = number that remain
n(knot) = starting number
gamma = (decay constant)
t = time
```
425
What is avogadros number
6.02 x 10^23
