Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

CHM135 > Thermochemistry > Flashcards

Thermochemistry Flashcards

1
Q

Define thermodynamics

A

the study of energy and its conversion from one form to another

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Define thermochemistry

A

a branch of thermodynamics, which highlights how heat is involved in chemical and physical transformations

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

there are two basic kinds of energy:

A

potential energy (Ep): stored energy
kinetic energy (Ek): the energy of motion

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Ek =

A

1/2mv^2

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

internal energy, E, of a system

A

the sum of all energy (potential and kinetic) of everything in that system

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

chemical bond energy

A

potential energy stored in molecular bonds

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

energy of molecular motion

A
  • translational, rotational, and vibrational
  • sometimes also called thermal energy
  • temperature is a common measure of the energy of molecular motion
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

define energy

A

the capacity to do work or supply heat

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

change in energy of a system =

A

work done on the system + heat flow in to the system
ΔE = w + q

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

energy is measured in

A

Joules or kiloJoules

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

1st law of thermodynamics

A

energy cannot be created or destroyed
ΔE(universe) = 0
the total energy of the universe is conserved

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

system

A

chosen to include whatever you are focusing on (and things that cannot usefully be separated from that)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

surroundings

A

everything not in the system

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

universe =

A

system + surroundings

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

how can we determine the energy change of a system of interest and why?

A

by measuring the energy change in the surroundings; if energy leaves the system, it must enter the surroundings (and vice versa)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

any change in the energy of a system must be accompanies by

A

an equal magnitude change in the energy of the surroundings, but the signs of these changes must be opposite

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

ΔE(universe) =

A

ΔE(system) + ΔE(surroundings) = 0

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

in thermodynamics, signs are defined from

A

the system’s point of view/perspective

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

ΔE of a reaction

A

= E(products) - E(reactants)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

if a reaction releases energy

A
  • system loses energy
  • reaction produces heat or does work
  • reactants higher in energy than the products
    ΔE < 0
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

state function

A

values that depend on the state of the substance, and not on how that state was reached. For example, density is a state function, because a substance’s density is not affected by how the substance is obtained.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

examples of state functions

A

E, V, P

23
Q

examples of non-state functions

A

q and w

24
Q

heat (q)

A

heat (q), also called thermal energy, will flow from higher-temperature objects to lower-temperature objects. eventually, the objects will reach the same temperature

25
Q

work (w)

A

mechanical work is the product of force (F) and distance (d)

w = F x d

the larger the required force, or the longer the distance on an object is moved, the more work is done on it

26
Q

example of energy transfer as work (w) only (q=0)

A

gas forming reaction in an insulated container that is attached to a piston-cylinder assembly that pushes against something
- the system pushes the piston out, doing work on the surroundings
- system releases (loses) energy
- ΔE = w < 0

27
Q

expansion work

A

the work done when the volume of a system changes in the presence of an external pressure

28
Q

expansion work is also often known as

A

pressure volume (PV) work because the amount of work done depends on both P and V

29
Q

at a constant pressure, P=P(surr), w=

A

-PΔV = -P(V(final) - V(initial))

more work is done when the volume change (ΔV) is larger and/or when pushing against a higher external pressure

30
Q

why is there a negative sign for expansion work?

A
  • required to fit with the convention that signs reflect the system’s perspective
  • when a system expands against a pressure:
    -> P, which is the constant pressure of the surrounding, is always >0
    -> the system does work. it loses internal energy by doing this work
    -> it must be that w<0, ΔE<0

we calculate: w = -PΔV <0

31
Q

enthalpy (ΔH)

A

defined as

H = E + PV
another accounting system for keeping track of energy, similar to internal energy, but does not include expansion work

32
Q

change in enthalpy (ΔH)

A

ΔH = ΔE + PΔV
- amount of heat absorbed or released in a transformation (real or imagined reaction) at a constant pressure.
- this means enthalpy change (ΔH) is often straightforward to determine via experiment

33
Q

the sign of ΔH shows

A

whether heat is produced or consumed in a reaction.

34
Q

ΔH<0

A

exothermic
- evolved heat flows out of the system into the surroundings
- heat is a product of the reaction

35
Q

ΔH>0

A

endothermic
- heat flows into the system from the surroundings
- heat is a reactant

36
Q

is enthalpy a state function?

A

yes.
ΔH = H(products) - H(reactants) = H(final) - H(initial)

37
Q

differences and similarities between enthalpy (ΔH) and internal energy change (ΔE)

A
  • differ in whether they account for the expansion (PV) work done on the surroundings or done by the surroundings on the system
  • most reactions involve little (if any) PV work. this is especially true for reactions involving only solids and liquids, or reactions in which the amount of gas does not change

thus, often (but not always) ΔH = ΔE

38
Q

effect of multiplying a reaction on ΔH

A

increases ΔH by the same factor

39
Q

effect of reversing a reaction on ΔH

A

changes the sign of ΔH for a reaction

40
Q

heat capacity

A

amount of heat (q, or ΔH) required to raise the temperature of an object or substance by 1’C (or 1K)

heat capacity = q/ΔT

41
Q

things with high heat capacity

A
  • require a lot of heat to increase in temperature and give off a lot of heat when they cool down.
  • act as a good sponge for (or, store of) thermal energy
42
Q

types of heat capacity

A

specific heat capacity (c): heat capacity per mass (1.00g)
c = q/(mass x ΔT)
units of J/gK

molar heat capacity (C or Cm): heat capacity per mole
C = q/(amount (mol) x ΔT)
units of J/molK

43
Q

effect of energy from heat, q

A

excites both translational motion of molecules and vibrations and rotations within and between molecules

44
Q

calorimetry

A

the science of measuring the heat exchanged in chemical reactions

45
Q

in calorimetry, the heat of reaction (qrxn) is measured

A

indirectly, by means of a calorimeter.
- if the reaction produces heat, the temperature of the surroundings increases and vice versa

46
Q

q(rxn)

A

= - q(calorimeter)
= -(q(vessel) + q(solution) + q(other))

47
Q

constant-pressure, or ‘coffee-cup’ calorimetry

A

q(system) = - q(calorimeter)
- reaction or physical transformation is done in an insulated container at a constant pressure in a bath (usually water) of known heat capacity
- heat generated thus tells us about the enthalpy of reaction

q = qp = ΔE + PΔV = ΔH

  • heat capacity of the calorimeter (often ~equal to that of the surrounding water because absorption of heat by the vessel is minimal) provides the link between change in temperature and heat gained or lost, from which you can determine the enthalpy change of the system

Heat capacity (calorimeter) = q(cal)/ΔT(cal)

48
Q

second type of calorimetry

A

constant volume or ‘bomb’ calorimetry:
- measures the heat change at constant volume such that q = qv = ΔE + PΔV = ΔE
- commonly used to measure heat of combustion reactions

49
Q

Hess’s Law

A

the enthalpy change for a process is equal to the sum of the changes for individual steps (1, 2, …n) of the process
ΔH(overall) = ΔH1 + ΔH2 + …. + ΔHn

50
Q

standard states

A
  • pure substance: most stable form at 1atm
  • gas: 1atm and ideal behaviour
  • substance in aqueous solution: 1M concentration all at a specified temperature, which is usually 25’C
51
Q

standard enthalpy of reaction

A

ΔH’(rxn)
- all reactants and products are in standard states

52
Q

ΔH’f

A

standard heat of formation
- enthalpy change for the formation of 1 mole of substance in its standard state from its constituent elements in their standard states

53
Q

we can use our database of ΔH’f’s to calculate ΔH (rxn)

A

ΔH’(rxn) = ΔH’f(products) - ΔH’f(reactants)
- ΔH of each reactant or product must be multiplied by its stoichiometric coefficient in the balanced rxn equation

54
Q

ΔH’(rxn) =

A

bonds broken - bonds made

CHM135 (23 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions