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IB Chemistry HL > Chapter 5/ 15 > Flashcards

Chapter 5/ 15 Flashcards

Energetics/thermochemistry

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1
Q

Heat

A

The transfer of energy between objects of different temperature

  • heat will spontaneous flow from an object of higher temp. to an object of lower temp.
  • once two objects reach same temp., which is known as thermal equilibrium, no more energy will be transferred
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2
Q

Enthalpy

A

Heat released or absorbed by a system at constant pressure

- is the heat content of a system

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3
Q

Standard enthalpy of reaction

A

When the enthalpy change is measured under standard conditions

  • pressure of 100 kPa
  • temp. of 25 degrees Celsius (298K)
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4
Q

Temperature

A

The average kinetic energy of the particles in a substance

  • if temp. of a substance increases, average kinetic energy of particles is increasing
  • absolute temp. is measured on Kelvin scale
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5
Q

Kelvin scale

A
  • absolute temp. (kelvin, K) is directly proportional to average kinetic energy of particles in a sample
  • absolute 0 (0K is lowest temp. theoretically possible), temp. at which all particles have least amount of kinetic energy
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6
Q

Kelvin - degrees Celsius equation

A

Temperature (K) = temperature (degrees Celsius) + 273.15

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7
Q

System and surroundings

A 
System= chemical reaction
Surroundings= rest of the universe
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8
Q

Open systems

A

Where matter and energy can move freely between system and surroundings

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9
Q

Closed systems

A

Only energy is able to move between system and surroundings

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10
Q

Relationship between system and surroundings

A
  • energy that is lost from system is transferred to surroundings
  • any energy gained by the system is transferred from surroundings to system
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11
Q

Isolated system

A
  • doesn’t exchange energy or matter with its surroundings
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12
Q

Negative change in enthalpy

A

When heat is released from the system to the surroundings, enthalpy of the system decreases

  • change in enthalpy is negative
  • only exothermic reactions
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13
Q

Positive change in enthalpy

A

When heat is transferred to a system from surroundings, enthalpy of system increases

  • change in enthalpy is positive
  • only endothermic reactions
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14
Q

Law of conservation of energy

A

Energy cannot be created or destroyed; it’s converted from one form to another
- hence, total amount of energy in universe is constant

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15
Q

Chemical reaction

A

Involves a transfer of heat between a system and its surroundings
- in majority of chemical reactions, heat is released from system to its surroundings (exothermic reactions)

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16
Q

Exothermic reactions

A
  • heat is released from system to its surroundings
  • in aqueous solutions, cause an increase in temp. of reaction mixture; it will feel hotter
  • warms the mixture, heat is lost to surroundings
    eg. combustion and neutralisation reactions
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17
Q

Energy profile for an exothermic reaction

A
  • enthalpy change is negative because heat has been transferred from system to surroundings
  • system has lost heat
  • products of reaction have lower enthalpy than reactants, so are more energetically stable
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18
Q

Thermochemical equations

A

Show enthalpy change of a reaction

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19
Q

Endothermic reactions

A

Chemical reactions in systems that absorb heat energy from their surroundings

  • in aqueous solutions cause a decrease in temp. of reaction mixture
  • it feels colder
  • reaction mixture is part of surroundings (not the system); system absorbs heat from its surroundings in an endothermic reaction
  • take in heat; cools the mixture at first, and then heat is gained from the surroundings
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20
Q

Energy profile for an endothermic reaction

A
  • enthalpy change is positive because reactants gain heat from their surroundings
  • products have higher enthalpy than reactants, and are less energetically stable
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21
Q

Coffee cup calorimeter

A
  • used to determine enthalpy changes
  • a polystyrene cup w/ a lid and a thermometer
  • reaction is carried out in the polystyrene cup, and temp. change of reaction mixture is recorded using thermometer
  • once min. and max. temp. of reaction mixture are known, enthalpy change of reaction can be calculated
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22
Q

Advantages of simple polystyrene cup as a calorimeter

A
  • polystyrene is a good heat insulator, it reduces heat loss to the surroundings
  • expanded polystyrene cup absorbs very little heat itself, no need to calculate heat absorbed by calorimeter
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23
Q

Calculating enthalpy changes

A

q = m x c x (change in temp.)

q = heat absorbed or released in in J
c = specific heat capacity in J/g/K
m = mass of solution in g
change in temp. = change in temp. in degrees Celsius or K

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24
Q

Specific heat capacity (c)

A

Amount of heat required to raise the temp. of one gram of a substance by 1 degree Celsius or 1 Kelvin

Units: J/g/degrees Celsius or J/g/K

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25
Q

Specific heat capacity of certain substances

  • metals
  • water
A

Metals:

  • lower specific heat capacity
  • heat up quickly, but also lose heat quickly

Water:

  • high specific heat capacity
  • takes a lot of heat energy to increase temp. of water
  • also retains that heat for a longer period of time

NB/ substances w/ higher specific heat capacities require more heat energy to increase their temp. and vice versa

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26
Q

Enthalpy change of neutralisation

A

Enthalpy change when an acid and base react together to form one mole of water

Method:

  • involves mixing known volumes and conc. of a strong acid and a strong base
  • measuring temp. increase
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27
Q

Method of measuring enthalpy change of neutralisation

A
  • a measured volume of a strong alkali is placed into polystyrene cup- equal volume of strong acid is added
  • temp. of reaction mixture increases until neutralisation is complete
  • continued addition of acid produces a cooling effect, no further reaction takes place
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28
Q

Standard enthalpy of combustion

A

The enthalpy change when one mole of a substance is burned completely in oxygen under standard conditions
- enthalpy changes of combustion are always negative- heat is released during combustion process

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29
Q

Method for standard enthalpy of combustion

A
  • known mass of alcohol is measured into a pre-weighed spirit burner and the alcohol is burned
  • heat released increases temp. of a known volume of water in calorimeter
  • temp. increase is measured for a certain time period, and experiment is then stopped
  • spirit burner and its contents are re-weighed
  • mass of alcohol burned to produce temp. increase is recorded, and molar enthalpy of combustion of alcohol can be calculated
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30
Q

Percentage error equation

A

% error = ((experimental -theoretical))/ theoretical) x 100

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31
Q

Limitations of calculating enthalpy changes in a school lab

A
  • heat loss to surroundings
  • incomplete combustion
  • assumptions made about specific heat capacity and density of aqueous solutions
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32
Q

State function

A

Value is independent of path taken to reach that specific value
- enthalpy (heat content of a system) is a state function

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33
Q

Hess’s law

A

States that, the total enthalpy change in a chemical reaction is independent of the route by which the chemical reaction takes place, as long as the initial and final conditions are the same

  • change in enthalpy during a chemical reaction doesn’t depend on whether reaction proceeds in one step or many steps
  • used to calculate enthalpy changes for reactions that can’t be determined experimentally
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34
Q

Enthalpy cycle

A
  • a cycle of reactants being converted into products
  • there are two routs, or paths, that a reaction can take

Direct route: conversion of reactants directly into products

Indirect route: conversion of reactants, via certain reaction intermediates, into products

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35
Q

Standard enthalpy change of formation

A

Enthalpy change when one mole of a compound is formed from the elements in their standards states under standard conditions

  • values indicate stability of compounds in relation to their elements
  • more -ve the value, greater the stability of the compound

Standard enthalpy of formation of an element in its standard state is 0

  • because, enthalpy change for formation of one mole of an element in its standard state from itself would be 0, as reactants and products would be the same
  • hence, there wouldn’t be a change in enthalpy
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36
Q

Chemical reaction

A

Involves breaking of existing chemical bonds and making of new chemical bonds

  • chemical bonds are results of electrostatic attractions between atoms or ions
  • to overcome these attractions (to break the bond), energy must be absorbed
  • bond breaking is endothermic (requires energy)

NB/ when new bonds are formed= energy is released
- bond making is exothermic

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37
Q

When is a chemical reaction exothermic?

A

In a chemical reaction, when energy absorbed while breaking bonds is less than energy released when forming new bonds

38
Q

When is a chemical reaction endothermic?

A

In a chemical reaction, when energy absorbed while breaking bonds is more than energy released when forming new bonds

39
Q

Overall enthalpy change for a reaction

A

Difference between energy absorbed and energy released

40
Q

Bond enthalpy (H)

A

Energy required to break one mole of chemical bonds in the gaseous state

  • Energy required to break a chemical bond
  • also called bond dissociation energy, E
  • bond enthalpy values are always positive- refer to bonds being broken (endothermic process)
  • when new bonds are formed, energy released is equal to energy that was absorbed to break the bond
41
Q

Average bond enthalpy

A

Enthalpy change when one mole of bonds are broken in the gaseous state for the same bond in similar compounds

  • same type of chemical bond in different compounds has different bond enthalpy values
  • identical bonds in molecules w/ two (or more) types of bond also have different bond enthalpy values
42
Q

Experimental vs. average bond enthalpy values

A

Enthalpy changes for reactions calculated from average bond enthalpy data may differ from experimental values
- because average values taken from same bonds in a range of similar compounds

43
Q

Free radical

A

A highly reactive species due to presence of an unpaired electron
- unpaired electron is represented by a dot

44
Q

Earth’s atmosphere

A
  • earth has an oxygen-rich atmosphere
  • atmosphere contains 2 forms of oxygen: di-oxygen and tri-oxygen (ozone)
  • both of these allotropes help protect life on Earth’s surface and in lower atmosphere by absorbing UV from the sun
45
Q

Early atmosphere of the earth

A
  • lacked any significant level of oxygen- gave little protection against UV radiation
  • this could have been a possible factor in rapid evolution, through mutation, of early life-forms
  • but following appearance of photosynthesising organisms, levels of atmospheric O2 increased
  • hence, came production of a protective, UV-absorbing layer of ozone
46
Q

Ozone layer

A

Stratosphere = 12-50km above surface of the Earth

  • lower regions contain 90% ozone in atmosphere; conc. less than 10ppm
  • stratospheric ozone is in a dynamic equilibrium w/ oxygen and is continually being formed and destroyed
  • ozone levels are maintained naturally by a continuous cycle of synthesis and breakdown reactions by action of energy from UV-C and UV-B radiation
47
Q

Ozone bonds vs. oxygen bonds

A
  • ozone bonds are weaker to those in oxygen (hence, results in phenomena of ozone generation and depletion
  • strong covalent bond in di-oxygen can be broken in the stratosphere (10-50km above Earth’s surface) by UV w/ wavelength shorter than 242nm, producing 2 free oxygen radicals
48
Q

Ozone generation

A

Bonds in di-oxygen are broken in stratosphere by UV radiation (wavelength shorter than 242nm) producing 2 free oxygen atoms

molecular oxygen —> atomic oxygen

Oxygen atoms produced each have an unpaired electron- highly reactive free radicals, react w/ oxygen molecules to form ozone

Reaction removes high-energy UV radiation, prevents it from reaching Earth’s surface
- results in this level of stratosphere having a higher temp. than lower region

49
Q

Ozone depletion

A

Bonds in ozone molecules are weaker than double bonds in di-oxygen- broken by UV radiation of lower energy

  • 2 reactions in this cycle depend on different wavelengths of UV radiation- effect is to remove higher energy radiation
  • only longer wavelength, less damaging radiation, reaches Earth’s surface
  • breakdown of ozone requires UV radiation of shorter wavelength than breakdown of ozone
  • strong double bond in oxygen is disrupted by Sun’s high energy UV-C radiation to form free radicals
  • these oxygen radicals can then react w/ an oxygen molecule to form ozone
  • bonds in ozone, being weaker, can then be broken by less energetic UV-B radiation (of longer wavelength) to reform oxygen and an oxygen free radical
50
Q

UV radiation and the earth

A

Surface of the earth is protected by Ozone depletion and ozone generation from damaging effects of UV-B and UV-C radiation

  • hence, majority of UV radiation reaching Earth’s surface is the least harmful UV-A form
  • ozone layer protects life on Earth from that radiation that would be most harmful to living tissues
51
Q

Ionic compound

A

Consists of oppositely charged ions held together in a lattice structure by strong electrostatic attractions

  • requires energy to overcome electrostatic attractions between ions
  • hence, bond breaking is endothermic
52
Q

Lattice enthalpy (lattice dissociation enthalpy)

A

The enthalpy change when one mole of a solid ionic compound breaks down to form gaseous ions under standard conditions

  • amount of energy required to separate an ionic compound into its gaseous ions
  • it is an endothermic process

NB/ It’s directly proportional to product of ionic charges
- inversely proportional to distance between nuclei of ions

53
Q

What is the opposite of lattice enthalpy?

A

Lattice formation enthalpy

54
Q

Lattice formation enthalpy

A

The enthalpy change when one mole of an ionic solid is formed from gaseous ions
- would have same numerical value as lattice dissociation enthalpy, but opposite sign

55
Q

Magnitude of lattice enthalpy of an ionic compound

A
  • a measure of the strength of the ionic bonds between the ions
  • greater the value of the lattice enthalpy, more energy is required to overcome attraction between ions and stronger the bond
56
Q

Magnitude of lattice enthalpy depends on:

A
  1. Charge of the ions (ionic charge)
    - charge on ions has significant effect on electrostatic attraction between ions and magnitude of lattice enthalpy
    - higher ionic charge = stronger electrostatic attraction between ions, leading to higher value of lattice enthalpy
  2. Size of the ions (ionic radii)
    - lattice enthalpy decreases as ionic radius of halide ion increases
    - weaker electrostatic attractions between ions as ionic radius increases

NB/

  • ionic charge increases + ionic radii decrease = lattice enthalpy increases
  • ionic charge decreases + ionic radii increase = lattice enthalpy decreases
57
Q

Born-Haber cycle

A

An energy cycle that can be used to calculate the lattice enthalpy of an ionic compound

58
Q

Steps of the Born-Haber cycle

A
  1. Enthalpy of formation- enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions
  2. Enthalpy of atomisation- enthalpy change when 1 mole of gaseous atoms is formed from an element in its standard state
  3. Bond dissociation energy: energy required to break 1 mole of bonds in its gaseous state
  4. Ionisation energy: energy required to remove 1 mole of electrons from 1 mole of gaseous atoms
  5. Electron affinity: energy released when 1 mole of electrons are added to 1 mole of gaseous atoms
59
Q

Stable ionic compounds

A
  • large negative values for their enthalpy of formation

- as more energy is released, compound becomes more stable

60
Q

Process of dissolution of ionic compounds

A

Process of dissolving ionic compounds

  1. Break-up of lattice structure; ionic solid is converted to gaseous ions
    - enthalpy change for this process = lattice enthalpy
  2. Once lattice structure has been broken down, separated gaseous ions are hydrated by surrounding water molecules
  3. During hydration, ion-dipole forces are formed between gaseous ions and dipoles of water molecules
  4. -ve dipoles are attracted to positively charged ions and +ve dipoles are attracted to negatively charged ions
61
Q

Enthalpy of hydration

A

Enthalpy change when one mole of gaseous ions dissolves in water to form a solution of infinite dilution

  • change in enthalpy that occurs when an ion becomes hydrated
  • formation of strong ion-dipole forces releases energy
  • so, enthalpy of hydration values are always exothermic

General equation:
X+(g) –> X+(aq)

62
Q

A solution of infinite dilution

A

A solution that has a large excess of water

- addition of more water wouldn’t cause any more heat to be released or absorbed

63
Q

Enthalpy change of solution

A

Enthalpy change when 1 mole of solute dissolves to form a solution of infinite dilution

64
Q

Solubility of an ionic compound depends on:

A
  • sign of enthalpy change of solution

- magnitude of enthalpy change of solution

65
Q

Insoluble ionic compounds

A

Relatively high positive value for enthalpy of solution (endothermic)

  • such values cause salt to be insoluble because combined enthalpy of hydration values of ions don’t compensate for high lattice enthalpy value
  • means enthalpy of solution value is endothermic, dissolving process is energetically favourable
66
Q

Soluble ionic compounds

A
  • negative exothermic value for enthalpy of solution
  • high enthalpy of hydration values compensation for endothermic enthalpy of lattice
  • results in an exothermic enthalpy of solution value
  • dissolving process is energetically favourable

Substances w/ slight positive (endothermic) values for enthalpy of solution can also be soluble

67
Q

Factors affecting enthalpy of hydration

A
  1. Ionic radii
    - as ionic radii increase, magnitude of enthalpy of hydration values decreases
    - due to weaker ion-dipole forces produced as size of ion increases
  2. Charge on ion (charge density of the ion)
    - as ionic charge increases, enthalpy of hydration values become increasingly exothermic
    - due to strength of ion-dipole force; higher the charge on the ion, stronger the force, more energy is released

Charge density of an ion increases w/ greater ionic charge and a decrease in ionic radius

  • higher charge density results in a stronger ion-dipole force between ion and water molecule
  • hence a greater, more negative value, for enthalpy of hydration
68
Q

Spontaneous reactions

A

For a reaction to be spontaneous:

  • must be an associated increase in total entropy of system together w/ surroundings
  • spontaneous change always increases total entropy of universe (2nd law of thermodynamics)
  • as entropy increases, amount of energy in universe available in a form to do useful work decreases
69
Q

Entropy (S)

A

Refers to distribution of available energy among particles in a system

  • more ways there is for energy to be dispersed, or spread out, greater the entropy
  • in a spontaneous process, energy goes from a concentrated form to a dispersed form
70
Q

What happens in a spontaneous process?

A

Occurs without addition of energy, other than that required to overcome initial energy barrier

71
Q

Combustion of coal

A
  • example of an increase in entropy for a spontaneous process
  • coal is a conc. source of energy that is burned in power stations to generate electricity
  • heat produced by combustion of coal heats water to produce steam, this is then used to drive a turbine and ultimately a generator
  • steam then passes to cooling towers of power stations where it condenses to form water
  • heat released when steam condenses is transferred to surrounding environment

Energy has gone from a conc. form in coal to dispersed form in surroundings

  • once in dispersed form, energy is less able to ‘do work’ = low-quality energy
  • energy’s natural tendency is to spread out and become less ‘useful’ over time
72
Q

Relative entropy of solids, liquids and gases

A

Solids:

  • fixed arrangements of particles
  • lowest entropy

Gases:

  • random arrangement of particles
  • highest entropy, because energy can be distributed in more ways than in liquids and solids

Entropy of gas > liquid > solid under the same conditions

73
Q

Entropy change (ΔS)

A
  • can be negative (−ΔS) or positive (+ΔS)
  • in chemical reactions that involve solids or liquids being converted into gaseous products, there will be a large increase in entropy
  • if there’s a decrease in no. of moles of gas on the product side, decrease in entropy
74
Q

Gibbs free energy (G)

A

relates the energy that can be obtained from a chemical reaction to the change in enthalpy (ΔH), change in entropy (ΔS) and absolute temperature (T)

75
Q

The change in Gibbs free energy (ΔG)

A

a convenient way to take into account both the direct entropy change resulting from the transformation of the chemicals, and the indirect entropy change of the surroundings due to the gain/loss of heat energy

76
Q

Standard entropy

A

Standard entropy of a substance is the consequential entropy change from heating the substance from absolute 0 (0K) to the thermodynamic standard temperature, 298K

  • a perfect crystal at 0K has a standard entropy of 0, hence, every other substance has a +ve standard entropy value
  • standard entropy values can be used to to calculate standard entropy change (ΔS⦵)
77
Q

Standard entropy change (ΔS⦵)

A

ΔS⦵ = ΣΔS⦵ (products) − ΣΔS⦵ (reactants)

78
Q

Spontaneous process

A

occurs without the addition of energy, other than that required to overcome the initial energy barrier

79
Q

Non-spontaneous reaction

A

one that needs a constant input of energy to occur

80
Q

Second law of thermodynamics

A

states that the total entropy of the universe tends to increase

  • tells us the direction of entropy change that favours a spontaneous chemical reaction or physical change taking place
  • for a spontaneous process, the total entropy of the system and surroundings (ΔStotal) must increase
81
Q

Total entropy of the system and the surroundings

A

ΔS total = ΔS system + ΔS surroundings ≥ 0

82
Q

Why do some chemical reactions have a decrease in entropy but are still spontaneous?

A
  • because the entropy of the surroundings increases to a much greater extent
  • this gives an overall increase in entropy for the process
  • increase in entropy of surroundings is usually as heat released from the system into the surroundings
  • depending on temp. of the surroundings, heat released can have a small or a large effect on the entropy of the surroundings
  • If the surroundings are cool, heat will have a large effect, but if surroundings are hot, heat will have a small effect
83
Q

Gibbs free energy change, G equation

A

ΔG⊖=ΔH⊖−TΔS⊖

84
Q

Gibbs free energy

A

The energy associated w/ a process that can be used to do work, specifically non-expansion work

85
Q

Work

A

Work done by a system can be:

  1. Expansion work
    - due to the change in the volume of a system by expansion or contraction
  2. Non-expansion work
    - work that isn’t associated w/ expansion of a system
    - ‘useful work’
86
Q

Free energy changes in a reaction

A
  • As a reaction proceeds, composition of the reaction mixture is constantly changing, as is the free energy
  • position of equilibrium corresponds to a max. value of entropy and a min. value of Gibbs free energy
  • once this point has been reached, reaction won’t proceed any further – rates of the forward reaction and reverse reaction are now equal
  • composition of an equilibrium mixture thus depends on value of ΔG⦵
87
Q

3 factors to consider when determining spontaneity of a reaction

A
  • the sign of the ΔH
  • the sign of the ΔS
  • temp. at which reaction takes place

Thus, ΔG = ΔH – TΔS

ΔH = enthalpy change in kJ/mol

T = absolute temp. in K (kelvin)

ΔS = entropy change in J/K/mol

88
Q

Change in enthalpy vs. change in entropy

A

ΔH units = kJ/mol

ΔS units = J/K/mol

  • necessary because unit for Gibbs free energy changes is kJ/mol
89
Q

Determining whether a reaction is spontaneous or not

A
  1. If Gibbs free energy change, ΔG, is -ve, reaction is spontaneous
  2. If Gibbs free energy change, ΔG, is +ve, reaction is non-spontaneous
  3. If Gibbs free energy change, ΔG, is 0, reaction is at equilibrium

NB/ sign of the ΔG value indicates spontaneity of a reaction
- it gives no indication of the rate at which the reaction takes place

90
Q

Exothermic and endothermic reactions and spontaneity

A
  1. Exothermic reactions w/ increase in entropy (-ve ΔH and +ve ΔS) are spontaneous regardless of temp.
  2. Endothermic reactions w/ decrease in entropy (+ve ΔH and negative ΔS) are non-spontaneous at any temp.

In both of these cases, signs of change in enthalpy and change in entropy mean that ΔG is independent of temp.

  1. Exothermic reactions w/ decrease in entropy are only spontaneous at low temp.
    – when product of TΔS is small enough to give a -ve ΔG value when subtracted from -ve ΔH value
    - At higher temp., product of TΔS, when subtracted from -ve ΔH value, gives a +ve ΔG value (non-spontaneous)
  2. Endothermic reactions w/ increase in entropy are only spontaneous at high temp.
    - If temp. is high enough, subtracting product of TΔS from ΔH value will give a -ve (spontaneous) ΔG value
    - at lower temp., product of TΔS isn’t large enough to cancel out +ve ΔH value, so the ΔG value is +ve (non-spontaneous)
91
Q

Gibbs free energy of formation (ΔG⦵f)

A

the change in free energy when one mole of a compound is formed from its elements in their standard states under standard conditions
- can be used to calculate standard Gibbs free energy change, ΔG⦵, for a reaction

92
Q

Standard Gibbs free energy change, ΔG⦵

A
  • the change in Gibbs free energy for a reaction at a temperature of 298 K

Equation:
ΔG⦵ = ΣΔG⦵f (products) − ΣΔG⦵f (reactants)

Results:

  • sign of ​ΔG⦵ value indicates stability of a compound relative to its constituent elements
  • Compounds w/ a -ve free energy change are more thermodynamically stable than their constituent elements
  • compounds w/ a +ve free energy change are less thermodynamically stable than their constituent elements

IB Chemistry HL (9 decks)

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