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chemistry 1A > thermodynamics introduction > Flashcards

thermodynamics introduction Flashcards

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1
Q

open system

A

an unsealed flask
there will be transfer of matter (molecules) and transfer of energy between the system and surroundings.

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2
Q

closed system

A

a sealed flask
there will be transfer of energy between the system and surrounding but no transfer of matter.

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3
Q

isolated system

A

a sealed flask in a vacuum flask
there will be no transfer of energy or matter between the reaction system and surroundings.

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4
Q

reaction equations note

A

when writing down a reaction equation at a specific temperature, the products must be in their specific state at that temperature, even if it produces the product in a different state, since the product will then change state to its normal state at that temperature.

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5
Q

1st law of thermodynamics

A

the internal energy of an isolated system is always constant.

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6
Q

2nd law of thermodynamics

A

the entropy of a reaction system and its surroundings is always increasing for a spontaneous process.

This basically means that a reaction will proceed spontaneously if the total entropy of the universe will be increased by the reaction.

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7
Q

3rd law of thermodynamics

A

the entropy of a perfect crystalline material at 0K = 0

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8
Q

1st law helpful development

A

the change in internal energy of a closed system is equal to the heat and work done by the system.

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9
Q

internal energy equation

A

ΔU = q + w
change in internal energy = heat energy + work

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10
Q

what is the potential energy of a reaction system

A

the stored energy of chemical bonds in the reaction system.

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11
Q

how do we calculate the potential energy of the reaction system

A

we use the mean molar bond enthalpies,

bonds broken are endothermic and bonds formed are exothermic, so do reactants - products.

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12
Q

mean molar bond enthalpy definition

A

the enthalpy change when a specific bond is broken within a molecule.

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13
Q

what is the kinetic energy of a molecule determined by

A

the number of degrees of freedom that the molecule has.

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14
Q

degrees of freedom definition

A

distinct ways in which molecules can move.

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15
Q

3 types of degrees of freedom

A

rotational
vibrational
translational

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16
Q

translational degrees of freedom

A

translational degrees of freedom are movements of a molecule through space in the x, y, and z axis.
there are always 3

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17
Q

rotational degrees of freedom

A

the degrees of freedom produced by bond rotations, there will be 3 in non linear molecules and 2 in linear molecules.

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18
Q

vibrational degrees of freedom

A

the degrees of freedom produced by the molecules bonds absorbing energy to bend and stretch. (normal modes)
3N - 5 for linear molecules
3N - 6 for non linear molecules.

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19
Q

kinetic energy per degree of freedom per molecule formula

A

E = 1/2 kb x T
energy = 1/2 Boltzmann’s constant x Temperature

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20
Q

Boltzmann’s constant

A

1.381 x 10^-23 J K-1

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21
Q

energy per degree of freedom per mole formula

A

E = 1/2 RT
energy = 1/2 gas constant x temperature

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22
Q

gas constant

A

8.314 J K-1 mol-1

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23
Q

kinetic and potential energy comparison

A

kinetic energy is in the 10’s of kJ and potential energy is in the 100’s, to potential > kinetic.

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24
Q

most important type of work in system

A

work of expansion - since it will have the largest value, so will contribute to most of the work.

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25
Q

work of expansion requirement

A

there must be a change in the number of moles of a gas in a reaction

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26
Q

properties of gasses (basic)

A

p ∝ 1/V at constant temperature
V ∝ T at constant pressure
p ∝ T at constant volume

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27
Q

how should you understand constant pressure and volume

A

constant pressure is like a piston pushing down on a chamber at consistent force
constant volume is like an indestructible box with something inside of it.

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28
Q

ideal gas law

A

pV = nRT
pressure x volume = number of moles of gas x gas constant x temperature.

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29
Q

ideal gas law assumptions

A

the gas particles are infinitely small points
the gas particles move in straight lines
the gas particles have no intermolecular forces of interaction between each other

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30
Q

mixture of gasses pressure formula
(Daltons law)

A

p = p1 + p2 + p3 +…
the pressure of a mixture of gasses = the pressure of all of the different components of the gas mixture added together.

where p1 = x1p

where x1 is the percentage composition of the mixture of gasses.

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31
Q

work of expansion formula

A

w = -p(ΔV)
work of expansion = minus pressure x change in volume

this can then be expanded using the ideal gas law

w = -ΔnRT
work of expansion = minus change in moles x gas constant x temp

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32
Q

work of expansion + or -

A

work of expansion of a system will be -ve since w = -ΔnRT and Δn is +ve.

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33
Q

work of contraction + or -

A

work of contraction of a system will be +ve since w = -ΔnRT and Δn will be -ve.

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34
Q

specific heat capacity definition

A

the energy required for 1g of a substance to have its temperature raised by 1K

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35
Q

molar heat capacity definition

A

the energy required for one mole of a substance to have its temperature raised by 1K

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36
Q

specific heat capacity formula

A

Cs = q/mΔT

37
Q

molar heat capacity formula

A

Cp = q/nΔT

38
Q

enthalpy change formula relating to internal energy, pressure and volume

A

ΔH = ΔU + pΔV

39
Q

enthalpy at constant pressure

A

ΔH = ΔU + pΔV
pΔV = -w
ΔH = ΔU - w
ΔU = q + w
ΔH = q + w -w
ΔH = q

40
Q

enthalpy at constant volume

A

ΔH = ΔU + pΔV
pΔV = 0 at constant volume
ΔH = ΔU = qv at constant volume

41
Q

exothermic reactions definition

A

a chemical reaction where there is heat energy transferred from the reaction system to its surroundings, meaning that heat is released and there is a -ve enthalpy change.

42
Q

endothermic reaction definition

A

a chemical reaction where there is heat energy transferred from the surroundings to the reaction system meaning heat is absorbed and there will be a +ve enthalpy change.

43
Q

forming bonds type of reaction

A

forming bonds will mean electrons will go from a high energy to low energy state, which will mean that heat energy is transferred from the system to the surroundings, which will mean that it is an exothermic reactions and the enthalpy change will be -ve

44
Q

breaking bonds enthalpy and type of reaction

A

breaking bonds will mean electrons go from a low energy state to high energy state, which will result in heat energy being absorbed by the system from surroundings, which will mean a +ve enthalpy change and an endothermic reaction.

45
Q

Hess law equation

A

Σ v Δf H°(products) - Σ v Δf H°(react)
note the v is just for stoichiometry

46
Q

Kirchhoff’s law equation

A

Δr H (T2) = Δr H (T1) + Δr Cp ΔT
change in enthalpy at T2 = change in enthalpy at T1 + molar heat x change in temperature.

47
Q

rule for enthalpy relating to temperature

A

as the temperature increases the enthalpy of the products and reactants will increase, and the change in enthalpy will increase since the difference between products and reactants will increase.

48
Q

important assumptions of Kirchhoff’s law

A

there cant be a state change
the molar heat capacity is assumed to be constant.

49
Q

latent heat definition

A

the enthalpy change required to change a substance from one state to another.

50
Q

fusion

A

solid to liquid

51
Q

vaporization

A

liquid to gas

52
Q

sublimation

A

solid to gas

53
Q

freezing

A

liquid to solid

54
Q

condensation

A

gas to liquid

55
Q

deposition

A

gas to solid

56
Q

entropy colloquial definition

A

the measure of disorder within a system

57
Q

entropy better definition

A

the measure of how energy is dispersed between different arrangements in a system whilst the systems energy remains constant.

58
Q

arrangements

A

W

59
Q

rule of entropy and number of atoms

A

As the number of atoms increases the number of different arrangements that energy can be dispersed between will increase, which will result in an increase in entropy.

60
Q

rule of entropy and state change

A

As a substance goes from a solid to a liquid to a gas the intermolecular forces holding the molecules in place break apart allowing the molecules to move more freely, which will result in a greater number of arrangements for energy to be dispersed between, meaning the entropy of the system will increase.

61
Q

which is bigger: the change in entropy between solids and liquids, or liquids and gasses?

A

liquids to gasses will have a greater change in entropy.

62
Q

how does temperature affect entropy

A

temperature will increase the average kinetic energy if the molecules, which will mean there will be more arrangements available for energy to be dispersed into.

63
Q

molar entropy of a reaction equation

A

Δr S°m = Σ Δr S°m (products) - Σ Δr S°m (reactants)

molar entropy at standard conditions = products molar entropy - reactants molar entropy

64
Q

entropy related to heat energy and temperature

A

ΔS = q(rev)/T
change in entropy = heat energy transfer of a reversible process/ temperature

65
Q

entropy at a different temperature formula

A

Δr S°(T2) = Δr S°(T1) + Cp ln(T2/T1)

entropy at temperature 2 = entropy at temperature 1 + molar heat capacity x ln T2/T1

66
Q

reaction feasibility for entropy rule

A

a reaction will proceed spontaneously if there is a total increase in the entropy of the universe, meaning the reaction will always move to a more distributed configuration if available.

67
Q

total entropy formula

A

Stot = Ssys +Ssur

68
Q

Gibbs free energy formula

A

ΔrG° = ΔrH° - TΔrS°

69
Q

rule for reaction feasibility in terms of free energy

A

if ΔG<0 then a reaction is feasible, since ΔG = -StotT

70
Q

reaction spontaneity ΔH < 0 ΔS <0

A

the entropy change is -ve, so the entropy of the reaction system will be decreasing, the enthalpy is -ve so the reaction surroundings will be absorbing heat energy and the entropy of the surroundings will be increasing. This will mean that the reaction spontaneity will be temperature dependent.

71
Q

reaction spontaneity at ΔH <0 ΔS >0

A

the entropy change will be positive, meaning that the reaction systems entropy will be increasing, and the reaction is exothermic which will mean that the surroundings entropy will be increasing, therefore the total entropy will always be greater than 0, so the reaction will always be feasible.

72
Q

reaction spontaneity at ΔH>0 ΔS< 0

A

the entropy change is -ve which means the entropy of the system will be decreasing, the reactions enthalpy is endothermic which will mean that the surroundings will lose heat energy and therefore lose entropy, this will result in the reaction system never being feasible since the entropy of the reaction will always be decreasing.

73
Q

reaction spontaneity at ΔH> 0 ΔS>0

A

the reaction systems entropy will be increasing because the change in entropy is +ve, however the enthalpy change is -
+ve meaning heat energy is being absorbed by the reaction system, which means the energy and entropy of the reaction surroundings is decreasing, meaning that the temperature will affect the reaction feasibility.

74
Q

formula for Δ G at standard condition

A

Δr G° = Σ v Δr G ° (products) - Σ v Δr G ° ( reactants)

75
Q

what does the reaction quotient describe

A

the ratio of products to reactants

76
Q

when does the reaction quotient equal the equilibrium constant

A

at equilibrium when the Gibbs free energy is at its minimum

77
Q

reaction quotient equation

A

Q = [C]^c [D]^d
[A]^a [B]^b

78
Q

equilibrium equation

A

K = [C]^c [D]^d
[A]^a [B]^b

79
Q

requirements for equilibrium

A

a reaction must not proceed to completion
there must be a minimum free energy within the mixture where there is not 100% products or reactants.
the reaction will always proceed towards this minimum free energy
once this minimum free energy is achieved the reaction rate of the forwards and reverse reaction will be constant, meaning that the concentrations of products and reactants will remain constant.

80
Q

why does a reaction not proceed to completion in equilibrium

A

because the minimum free energy in the reaction occurs when there are products and reactants present, and since ΔG will always be negative going towards the minimum free energy, the reaction will proceed until this minimum free energy, then products and reactants concentrations will remain constant.

81
Q

why do reactants and products remain constant at the minimum free energy in equilibrium

A

the reaction of the forwards and reverse reaction will occur at the exact same rate, so the reaction concentrations will remain constant

82
Q

which way does a reversible reaction proceed in a relation to Gibbs free energy

A

towards the minimum in Gibbs free energy.

83
Q

free energy at non standard conditions

A

Δr G = Δr G° + RT lnQ

84
Q

free energy at standard states related to equilibrium (vant hoffs isotherem)

A

Δr G° = - RT lnK

since Δr G = 0 at equilibrium and Q = K

85
Q

calorimetry at constant volume pressure procedure

A

calibrate the calorimeter with hot and cold water
measure the temperature changes of the hot and cold water
use -CsMΔT2 = CsMΔT1 + C cal ΔT1
to find the calorimeter constant.
then use ΔH = Cs M ΔT + cal ΔT

86
Q

calorimeter constant equation

A

-CsMΔT2 = CsMΔT1 + C cal ΔT1
heat energy lost by hot water = heat energy gained by cold water + heat energy lost by calorimeter.

87
Q

enthalpy change of a calorimeter at constant pressure equation

A

ΔH = Cs M ΔT + cal ΔT

enthalpy change = specific heat capacity x mass x temperature change + calibration constant x temperature change

88
Q

calorimetry at constant volume explanation
(bomb calorimeter)

A

At constant volume no work of expansion can be lost, so the change in internal energy = the heat lost from the calorimeter chamber, which we can calculate using the calorimeter constant and the change in temperature.

89
Q

calorimetry at constant volume equation

A

ΔU = Ccal ΔT

chemistry 1A (9 decks)

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