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My Packs: (UWO) (CHEM 1302) > Chapter 2: 2.1 Heat, Work, and Energy > Flashcards

Chapter 2: 2.1 Heat, Work, and Energy Flashcards

1
Q

Define:

Thermodynamics

A

The study of the energy changes involved in physical and chemical processes

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2
Q

Define:

Thermochemistry

A

The branch of thermodynamics that investigates the heat flow that occurs during these reactions

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3
Q

Explain the difference between a “food/large calorie” (Cal) and a “small calorie” (cal)

A

Food calories are typically seen on food packages, and are actually 1000 cal
Small calories are used in thermodynamics

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4
Q

Define:

1 calorie

A

1 small calorie is equal to 4.184 Joules

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5
Q

Define

1 Calorie

A

1 Cal = 4.184 kJ

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6
Q

What is another well known definition of a large/food calorie?

A

Approximately the amount of energy needed to raise the temperature of 1 kg of water by 1 degrees Celsius (4.184 kJ)

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7
Q

True or False:

Heat always flows from cool to warm

A

False, heat always flows from the warmer object to the cooler object

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8
Q

State:

3 types of systems

A
  1. Open system
  2. Closed system
  3. Isolated system
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9
Q

Define:

Open system

A

Can exchange both matter and energy with its surroundings

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10
Q

Define:

Closed system

A

Can exchange energy but not matter with its surroundings

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11
Q

Define:

Isolated system

A

Exchanges neither matter nor energy with its surroundings

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12
Q

Define:

Energy

A

“The capacity to do work”
Work is defined as the product of a force acting on an object and the distance that the object moves in response to the force

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13
Q

What type of work are chemists usually interested in?

A

Pressure-Volume (PV) Work

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14
Q

Define:

Pressure-Volume Work

A

Work involved in the expansion or compression of gases

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15
Q

State and formula and units of:

PV work

A

w = -P * ΔV
Units in kPa L (1 kPa L = 1 J)

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16
Q

List and describe:

Some types of energy

A
  • Heat energy (thermal energy): The energy transferred due to a temperature difference bewteen the system and the surroundings
  • Kinetic energy: The energy associated with motion
  • Potential energy: Stored energy or the energy a body possesses due to its position
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17
Q

True or False:

Chemical energy is a form of thermal energy

A

False, chemical energy is a form of potential energy

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18
Q

True or False:

Energy can be created

A

False, energy cannot be created/destroyed; it can only be converted

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19
Q

Define:

Heat capacity

A

The parameter used to estimate the heat energy associated with temperature changes

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20
Q

What is a related quantity to heat capacity?

A

Specific heat capacity

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21
Q

What is the amount of heat transferred denoted by? What is it related to?

A
  • Denoted by q
  • Related to temperature difference (ΔT) and heat capacity/specific heat capacity
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22
Q

For heat capacity, state:

  • What it is used for
  • Notation
  • Units
A
  • Usually used for objects
  • Represented by the symbol C
  • Units: Joules/Degrees Celsius
23
Q

State the relation/formula for:

Heat capacity

(State the variables)

A

q = C * ΔT
* q: Amount of heat energy
* C: Object’s heat capacity
* ΔT: Temperature change

24
Q

For specific heat capacity, state:

  • Definition
  • Notation
  • Units
A
  • The heat needed to warm one mass unit (g or kg) of a substance by one degree
  • Represented by lowercase c
  • SI unit: Joules/Kilogram * Degrees Celsius
  • Common unit: Joules/Gram * Degrees Celsius
25
Q

State the relation/formula for:

Specific heat capacity

A

q = m * c * ΔT

26
Q

True or False:

When calculating amount of heat transferred, all Celsius units can be swapped with Kelvin units

A

True, as we are using the temperature difference

27
Q

Define:

Temperature change

A

The difference between final and initial temperatures
* ΔT = Tfinal - Tinitial

28
Q

When calculating temperature difference:

What do we have to ensure?

A

Tfinal and Tinitial are using the SAME SCALE

29
Q

When a system absorbs energy from the surroundings…

A

The process is endothermic and change in energy is positive

30
Q

When a system releases energy to the surroundings…

A

The process is exothermic and the change in energy is negative

31
Q

State the relation between:

ΔT and energy change (q)

A
  • -ΔT = -q, system cools
  • +ΔT = +q, system warms
32
Q

True or False:

The sign of q is the same as sign of ΔT

A

True

33
Q

Does heat move from cool to warm or the other way?

A

Heat always goes from warm to cool

34
Q

In heat transfer between objects:

State the change of heat, ΔT, and q for the warmer object

A
  • Loses heat and cools down
  • ΔT < 0
  • q (lost) is negative
35
Q

In heat transfer between objects:

State the change of heat, ΔT, and q for the cooler object

A
  • Gains heat and warms up
  • ΔT > 0
  • q (gained) is positive
36
Q

What must the number of joules lost by the warmer object equal to?

A

The number of joules gained by the cooler object

37
Q

What is it called when both objects reach the same final temperature (Tf)? What is a formula that represents this?

A

Thermal Equilibrium
* -q (lost) = q (gained)

38
Q

Define and give examples of:

Intensive (intrinsic) property

A

A property that has the same value regardless of the sample size
* E.x. Temperature (melting and boiling points), Density

39
Q

Define and give examples of:

Extensive (extrinsic) property

A

A physical property whose value increases with the sample size
* E.x. Mass, Volume, Energy, Enthalpy, and Entropy

40
Q

Define and give examples of:

State function

A

A property whose value depends ony on the current state of the system, not on how that state was reached
* E.x. Volume, Pressure, Enthalpy
* NOT Heat and Work

41
Q

State:

1st Law of Thermodynamics

A

The total energy of an isolated system is conserved

42
Q

How do we keep track of the energy of a system?

A

A thermodynamic function called * Internal Energy (E)*

43
Q

Define:

Internal energy (E)

A

E is the sum of all the kinetic and potential energies of all the atoms, ions, and molecules in the system

44
Q

Can internal energy be measured?

A

No it cannot, but change in internal energy can be measured

45
Q

What are some ways to measure ΔE?

A

q - heat
w - work

46
Q

For a system at constant pressure, how is work calculated?

A

w = -P * ΔV

47
Q

In gas expansion, state:

  • How is work being done
  • Value of w
  • Change in work energy
A
  • Gas (system) does work on surroundings
  • w is negative
  • Lost work energy
48
Q

In gas compression, state:

  • How is work being done
  • Value of w
  • Change in work energy
A
  • Surroudnings does work on gas (system)
  • w is positive
  • Gained work energy
49
Q

What happens if the process does not involve a compression or expansion? What would be the value of w?

A
  • The change in volume would be zero
  • The value of w is zero
50
Q

What does ΔE equal?

A

ΔE = q + w

51
Q

Determine what happens to the internal energy of the system when:

Surroundings do work on, or supply heat to the system

A
  • Internal energy of the system increases
  • ΔE > 0
52
Q

Determine what happens to the internal energy of the system when:

The system does work on, or supplies heat to the surroundings

A
  • Internal energy of the system decreases
  • ΔE < 0
53
Q

When the system is kept at a constant volume, what happens to the relationship with ΔE?

A
  • w = 0
  • The relationship becomes E = q, commonly written as E = qv

The subscript v in qv signifies the constant volume

54
Q

Theorem:

What does the internal energy changes of a system equal when the volume doesn’t change?

A

The internal energy change of a system is equal to the heat transferred when the volume does not change

My Packs: (UWO) (CHEM 1302) (15 decks)

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