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chemistry for engineers finals notes > GE-CHEM 1103 Module 4.1 > Flashcards

GE-CHEM 1103 Module 4.1 Flashcards

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1
Q

the ability to do
work or transfer heat

A

Energy

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2
Q

the study of energy and its
transformations

A

Thermodynamics

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3
Q

the study
of chemical reactions and the energy
changes that involve heat

A

Thermochemistry

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4
Q

most important form
of potential energy in
molecules

A

electrostatic
potential energy (Eel = kQ1Q2/d)

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5
Q

the unit of
energy commonly used

A

Joule ( 1J = 1 kg m^2/s^2)

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6
Q

it is
seen between
oppositely charged ions

A

Electrostatic attraction

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7
Q

what happens to energy when chemical bonds are formed?

A

Energy is released

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8
Q

what happens to energy when chemical bonds are broken?

A

Energy is
consumed

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9
Q

State the First Law of Thermodynamics

A

Energy can be converted from one form to
another, but it is neither created nor destroyed.

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10
Q

The portion of the
universe that we single
out to study is called

A

system

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11
Q

are
everything else that is not being studied

A

surroundings

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12
Q

a region of the
universe being studied that
can exchange heat AND
mass with its surroundings.

A

Open System

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13
Q

a region of
the universe being studied
that can ONLY exchange
heat with its surroundings

A

Closed System

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14
Q

a region of
the universe that can NOT
exchange heat or mass with
its surroundings

A

Isolated System:

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15
Q

the sum of all
kinetic and potential energies of all components
of the system

A

Internal energy; E
But we dont know E only how it changes so ΔE.

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16
Q

final energy of the system minus the initial
energy of the system

A

change in internal energy, ΔE
ΔE = Efinal − Einitial

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17
Q

the system absorbed energy from the
surroundings

A

ΔE > 0, Efinal > Einitial

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18
Q

the system released energy to the
surroundings

A

ΔE < 0, Efinal < Einitial

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19
Q

What is the value of ΔE if Efinal equals Einitial

A

ΔE = 0

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20
Q

Thermodynamic Quantities
Have Three Parts

A

1) A number
2) A unit
3) A sign

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21
Q

results when the system
gains energy from the surroundings.

A

positive ΔE

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22
Q

results when the system
loses energy to the surroundings.

A

negative ΔE

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23
Q

When energy is
exchanged between
the system and the
surroundings, it is
exchanged as either
heat (q) or work (w).

A

ΔE = q + w

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24
Q

Sign conventions for q

A

+q = system gains heat
-q = system loses heat

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25
Q

Sign conventions for w

A

+w = work done on system
-w = work done by system

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26
Q

Sign conventions for ΔE

A

+ΔE = net gain of energy by system
-ΔE = net loss of energy by system

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27
Q

When heat is absorbed by the system from the
surroundings

A

endothermic

28
Q

When heat is released by the system into the
surroundings

A

exothermic

29
Q

internal energy of a system

A

Is a state function (ΔE depends only on Einitial and Efinal)

30
Q

q(heat) and w(work)

A

Is not a state function

31
Q

the mechanical
work associated
with a change in
volume of gas

A

the only
work done by
chemical or
physical change

32
Q

measure the work done by the gas if the
reaction is done in a vessel that has been fitted with a piston. work is NEGATIVE because it is work done
BY the system

A

w = −PΔV

33
Q

the internal energy plus the
product of pressure and volume

A

Enthalpy

34
Q

When the system changes at constant
pressure, the change in enthalpy, ΔH, is

A

ΔH = ΔE + PΔV to
Since ΔE = q + w and w = −PΔV, we
can substitute these into the enthalpy
expression:
ΔH = ΔE + PΔV
ΔH = (q + w) − w
ΔH = q

35
Q

at constant pressure, the change in
enthalpy is

A

the heat gained or lost

36
Q

ΔH is
positive

A

process is
endothermic

37
Q

ΔH is negative

A

process is
exothermic

38
Q

is the
enthalpy of the
products minus the
enthalpy of the
reactants

A

The change in
enthalpy, ΔH

39
Q

ΔHrxn

A

enthalpy of
reaction, or the heat of reaction

40
Q

The Truth about Enthalpy of reaction

A
  1. Enthalpy of reaction is an extensive property.
  2. The enthalpy change for a reaction is
    equal in magnitude, but opposite in
    sign, to ΔH for the reverse reaction.
  3. The enthalpy change for a reaction
    depends on the states of the reactants
    and the products.
41
Q

is
defined as the enthalpy change for the
reaction in which a compound is made
from its constituent elements in their
elemental forms

A

enthalpy of formation, ΔHf

42
Q

are
measured under standard conditions (25 °C
and 1.00 atm pressure).

A

Standard enthalpies of formation, ΔHf
°

43
Q

enthalpy of formation for an element
in its elemental state

A

0 because it takes no energy to form a
naturally-occurring compound.

44
Q

Calculation of ΔH (see the answer right way)

A

Do some exercises on it

45
Q

Calculation of ΔH using Hess law

A

ΔH = ΣnΔHf,products – ΣmΔHf
°,reactants

46
Q

The enthalpy associated with breaking one
mole of a particular bond in a gaseous
substance

A

Bond Enthalpy

47
Q

always positive because energy is
required to break chemical bonds

A

bond enthalpy

48
Q

always released when a bond forms between
gaseous fragments

A

Energy

49
Q

greater the bond enthalpy

A

stronger the bond.

50
Q

ALL bonds made

A

ADD bond energy

51
Q

ALL bonds
broken

A

SUBTRACT bond energy

52
Q

Σ (bond enthalpies – Σ (bond enthalpies
of bonds broken) of bonds formed)

A

ΔHrxn

53
Q

If a reaction is carried out in a series of steps,
ΔH for the overall reaction equals the sum of the enthalpy
changes for the individual steps.

A

Hess’s law

54
Q

measure ΔH through measurement of heat flow

A

calorimetry

55
Q

instrument used to
measure heat flow

A

calorimeter

56
Q

The amount of energy required to raise the temperature of a
substance by 1 K (1 °C)

A

heat capacity

57
Q

the amount of
the substance heated is one gram

A

specific heat

58
Q

If
the amount is one mole

A

molar heat capacity

59
Q

The specific heat for water

A

4.184 J/g∙K (can be used for dilute solutions)

60
Q

calculate ΔH for the reaction

A

qsoln = Cs × msoln × ΔT = –qrxn

61
Q
  1. Reactions can be carried
    out in a sealed “bomb”
  2. Because the volume in the
    bomb calorimeter is
    constant, what is measured
    is really the change in
    internal energy, ΔE, not ΔH.
A

Bomb Calorimetry

62
Q

The heat absorbed (or
released) by the water is
a very good approximation
of the enthalpy change for
the reaction.

A

qrxn = – Ccal × ΔT

63
Q

The energy released when one gram of food is
combusted

A

fuel value

64
Q

Most of the energy in foods comes from carbohydrates,
fats, and proteins

A

Carbohydrates (17 kJ/g)
Fats (38 kJ/g)
Proteins (17 kJ/g)

65
Q

The vast majority of the
energy consumed in
this country

A

fossil fuels

66
Q

8.6% of the
U.S. energy needs

A

Nuclear fission

67
Q

produce 9.9% of the
U.S. energy needs

A

Renewable energy
sources

chemistry for engineers finals notes (6 decks)

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