CHEM 105 Ch. 9-10 Flashcards
(129 cards)
1
Q
energy
A
anything that has the capacity to do work
2
Q
work
A
a force acting over a distance (energy = work = force x distance)
3
Q
heat
A
the flow of energy caused by a difference in temperature
4
Q
kinetic energy
A
energy of motion or energy that is being transferred
5
Q
thermal energy
A
energy associated with temperature (a form of kinetic energy)
6
Q
potential energy
A
energy that is stored in an object or energy associated with the composition and position of the object
7
Q
energy stored in the structure of a compound is
A
potential energy
8
Q
chemical energy
A
potential energy due to the structure of the atoms, the attachment between atoms, the atoms’ positions relative to each other in the molecule, or the molecules’ relative positions in the structure
9
Q
nuclear energy
A
potential energy in the nucleus of atoms
10
Q
light/radiant energy
A
kinetic energy associated with energy transitions in an atom
11
Q
heat/thermal energy
A
kinetic energy associated with molecular motion
12
Q
electrical energy
A
kinetic energy associated with the flow of electrical charge
13
Q
the amount of kinetic energy an object has is directly proportional to
A
its mass and velocity (KE = 1/2 mv^2)
14
Q
when the mass is in kg and the velocity is in m/s, the unit for KE is
A
a joule (J) 1 J = kg*m^2/s^2 = 1 N*m
15
Q
one joule of energy is the amount of energy needed to move a ? kg mass at a speed of ? m/s
A
1, 1
16
Q
a calorie (cal) is
A
the amount of energy needed to raise the temperature of 1g water by 1*C
17
Q
1 kcal = energy needed to
A
raise 1000g water by 1*C
18
Q
a food Calorie (Cal) is ? calories
A
1000
1 Cal = 1 kcal
19
Q
1 calorie = ? joules
A
4.184
20
Q
1 kWh = ? joules
A
3.60 x 10^6
21
Q
thermodynamics
A
the study of energy that is exchanged between the system and surroundings
22
Q
the first law of thermodynamics
A
the law of conservation of energy
23
Q
the law of conservation of energy means that
A
the total amount of energy in the universe is constant; it is conserved (therefore, you can never design a system that will continue to produce energy without some source of energy)
24
Q
system
A
the part of the universe that is being studied
25
surroundings
everything else in the universe with which the system can exchange energy
26
when energy flows from the system to the surrounds,
the energy of the system decreases, the energy of the surroundings increases; exothermic
27
when energy flows from the surroundings to the system,
the energy of the system increases, the energy of the surroundings decreases; endothermic
28
conservation of energy means that the amount of energy gained or lost by the system has to be (less than/greater than/equal to) the amount of energy lost or gained by the surroundings
equal to
29
internal energy
the sum of the kinetic and potential energy of all the particles that compose the system
30
the change in the internal energy of a system only depends on
the amount of energy in the system at the beginning and end
31
state function
a mathematical function whose result only depends on the initial and final conditions, not on the process used
32
energy diagram
a "graphical" way of showing the direction of energy flow during a process
33
if the reactants have a lower internal energy than the products, then the change in energy will be
positive (endothermic)
34
if the reactants have a higher internal energy than the products, then the change in energy will be
negative (exothermic)
35
energy is exchanged between the system and surroundings through
heat and work
36
if q (heat) is positive
system gains thermal energy
37
if q (heat) is negative
system loses thermal energy
38
if w (work) is positive
work is done on the system
39
if w (work) is negative
work is done by the system
40
if ΔE (change in internal energy) is positive
energy flows into the system
41
if ΔE (change in internal energy) is negative
energy flows out of the system
42
heat is the exchange of
thermal energy between a system and surroundings
43
heat exchange occurs when
system and surroundings have a difference in temperature
44
temperature
the measure of the thermal energy within a sample of matter
45
heat flows from matter with (low/high) temperature to matter with (low/high) temperature until both objects reach the same temperature (thermal equilibrium)
high to low
46
increase in temperature is directly proportional to ? and the proportionality constant is called the ?
the amount of heat absorbed; heat capacity (C)
47
units for heat capacity (C) are ? and the equation is ?
J/*C or J/K
| q = cΔT
48
the larger the heat capacity of the object being studied, the (smaller/larger) the temperature rise will be for a given amount of heat
smaller
49
factors affecting heat capacity
amount of matter, type of material
50
the heat capacity of an object is proportional to
its mass and the specific heat of the material
51
the quantity of heat absorbed by an object can be determined if the following are known with the equation ?
mass, specific heat capacity, temperature change
| q = m x C x ΔT
52
specific heat capacity
measure of a substance's intrinsic ability to absorb heat; the amount of heat energy require to raise the temperature of 1g of a substance by 1*C; Cs; units J/(g*C)
53
the molar heat capacity is
the amount of heat required to raise the temperature of one mole of a substance by 1*C
54
? has the highest Cs
water (4.18 J/g*C)
55
pressure volume work is caused by
a volume change against an external pressure
56
when gases expand, ΔV is (negative/positive), but the system is doing work on the surroundings, so w,gas is (negative/positive)
positive, negative
57
as long as the external pressure is kept constance, w =
-PΔV
58
1 atm * L = ? J
101.3 J
59
exchange of heat energy equation, exchange of work equation
```
q = mCΔT
w = -PΔV
```
60
at constant volume, ΔEsystem =
q,system
61
in practice, temperature changes of individual chemicals involved in the reaction cannot be observed directly, so instead the ? in surroundings is measured using ?, so ?
ΔT, an insulated container, q,system = -q,surroundings
62
bomb calorimeter
used to measure ΔE because it is a constant volume system; the heat capacity of the calorimeter is the amount of heat absorbed by the calorimeter for each degree rise in temp. and is called the calorimeter constant
63
enthalpy, H, of a system
the sum of the internal energy of the system and the prouct of pressure and volume; a state function; H = E + PV
64
the enthalpy change, ΔH, of a reaction
the heat evolved in a reaction at constant pressure; ΔH,reaction = q,reaction at constant pressure
65
when ΔH is negative,
heat is being released by the system into the surroundings; exothermic
66
when ΔH is positive,
heat is being absorbed by the system from the surroundings; endothermic
67
the enthalpy of change in a chemical reaction is an ? property, meaning
extensive; dependent on the amount (i.e., more reactions => larger enthalpy change)
68
by convention, we calculate the enthalpy change for the number of moles of reactants in the reactions...
as written
69
ΔHreaction =
q,constant pressure = q,reaction
70
to get ΔHreaction per mole of a particular reactant,
divide by the number of moles that reacted then multiply by the coefficient in the equation
71
Hess's Law
the change in enthalpy for a step-wise process is the sum of the enthalpy changes of the steps
72
ΔHreaction can be estimated by comparing the cost of
breaking old bonds to the income from making new bonds
73
bond energy
the amount of energy it takes to break one mole of a bond in a compound
74
ΔHreaction = (in terms of bond energies)
Σ(ΔH(bonds broken)) + Σ(ΔH(bonds formed))
75
the more e-s two atoms share, the ? the covalent bond (must be comparing bonds b/w like atoms)
stronger
76
the shorter the covalent bond, the ? the covalent bond (must be comparing similar types of bonds) because...
stronger; bonds get weaker down the column and stronger across the period (atomic radii decrease across the period and increase down the column)
77
standard state
state of a material at a defined set of conditions
78
the standard enthalpy change, ΔH*, is the enthalpy change when
all reactants and products are in their standard states
79
standard state for gas
1 atm pressure
80
standard state for solid or liquid
1 atm pressure, usually 25*C temperature
81
standard state for substance in solution
concentration 1 M
82
the ΔH*,f for a pure element in its standard state
0 kJ/mol
83
ΔH*,rxn = (in terms of standard enthalpy)
Σ(nΔH*,f(products)) - Σ(nΔH*,f(reactants))
| where n is the coefficient for each product or reactant
84
crystal lattice
structure formed by an ionic solid in which every cation is surrounded by anions and vice versa; maximized the attractions b/w cations and anions, leading to the most stable arrangement; held together by the electrostatic attraction of the cations for all the surrounding anions
85
lattice energy
the extra stability that accompanies the formation of the crystal lattice; the energy released when the solid crystal forms from separate ions in the gas state (always exothermic); depends directly on the size of charges and inversely on distance b/w ions
86
Born-Haber cycle
a hypothetical series of reactions that represents the formation of an ionic compound from its constituent elements
87
the reactions in the Born-Haber cycle are chosen so that
the change in enthalpy of each reaction is known except for the last one, which is the lattice energy
88
how to use the Born-Haber cycle and Hess's law to calculate lattice energy
use Hess's law to add up enthalpy changes of other reactions to determine lattice energy
89
ΔH*(crystal lattice) =
lattice energy
90
ΔH*(metal atom (g)) =
first ionization energy
91
ΔH*(nonmetal atom (g)) =
electron affinity
92
trends in lattice energy
ion size and ion charge (more important)
93
the force of attraction between charged particles is (directly/inversely) proportional to the distance between them
directly (this is Coulomb's Law: q1q2/r)
94
a gaseous atom or molecule exerts a force when it collides with
a surface or other gaseous particles (molecular collisions are pressure)
95
gas pressure
the force exerted per unit area by gas molecules as they strike surfaces around them (P = F/A); a result of the constant movement of the gas molecules and their collisions with the surfaces around them
96
the pressure of a gas depends on
the number of gas particles in a given volume, the volume of the container, and the average speed of the gas particles
97
according to the Kinetic Theory of Gases, collisions of gas particles with each other and/or surfaces are said to be ?, meaning ?
elastic, meaning no exchange of energy occurs
98
pressure (decreases/increases) with a higher concentration of gas molecules
increases
99
pressure is (low/high) with a low number of gas particles in a given volume
low
100
as volume increases, the concentration of gas molecules (decreases/increases)
decreases
101
1 mmHg = ? torr
1 torr
102
1 atm = ? mmHg = ? torr
760
103
1 atm = ? bar
1.013
104
1 atm = ? Pa
101325
105
1 atm = ? Psi
14.7
106
1 atm = ? in Hg
29.92
107
how manometers work
the difference in liquid levels is a measure of the difference in pressure between gas and atmosphere (if atmosphere side higher, gas has higher pressure b/c pushing liquid that way)
108
4 basic properties of gases and their units
1. pressure (atm)
2. volume (L)
3. temperature (K)
4. amount in moles (n)
109
Boyle's Law
pressure and volume are inversely related (P1V1 = P2V2)
110
Charles's Law
volume and temperature are directly related (V1/T1 = V2/T2)
111
Avogadro's Law
volume is directly proportional to the number of gas molecules (V1/n1 = V2/n2)
112
Ideal Gas Law
```
PV = nRT, where
P = pressure in atm
V = volume in L
n = # moles
R = constant (0.0821 L*atm/mol*K)
T = temperature in K
```
113
standard conditions
```
P = 1 atm
T = 273K = 0*C
n = 1 mol
V = 22.4 L
```
114
density of a gas
molar mass of gas (g/mol) / molar volume (L/mol)
| molar volume = 22.4 L
115
molar mass of a gas
mass (g) / mole (n)
116
gas density from PV=nRT
```
d = PM / RT
(M = molar mass; can use n = mass (g) / M)
```
117
dry air is composed of
nitrogen, oxygen, argon, carbon dioxide, and a few other gases
118
Dalton's Law
total pressure of a gas mixture = sum of partial pressures
119
particles with different masses have (the same/different) kinetic energies at a given temperature
the same
120
partial pressure of a gas
the pressure of a single gas in a mixture of gases
121
partial pressure can be calculated if
1) a fraction of the mixture it composes and the total pressure are known
2) the number of moles of the gas in a container of a given volume and temperature is known
122
partial pressure equation
Pn = n,n * (R*T / V) or X,a * P,total
123
total pressure from partial pressures equation
P,total = n,total * (R*T / V)
124
mole fraction, X,a
the ratio of the partial pressure a single gas contributes to the total pressure; X,a = P,a / P,total = n,a / n,total
125
deep sea diving and oxygen levels
abnormally high partial pressure of O2 => elevated concentration of oxygen in body tissues; safe range is between 0.21 and 1.4
126
collecting gases by displacing water from a container
get the gas + water vapor (the partial pressure of the water vapor depends on temperature only)
127
Kinetic Molecular Theory
the simplest model for the behavior of gases; a gas is modeled as a collection of particles in constant motion; the average KE of the gas particles is directly proportional to the Kelvin temperature; collisions completely elastic
128
postulates of the KMT
the particles of the gas are constantly moving, the attraction b/w particles is negligible, collisions => bounce off & keep moving, lots of empty space b/w gas particles compared to size of particles
129
chapter ten
continued
