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03 Chemistry > Module 4 (Drivers of Reactions) > Flashcards

Module 4 (Drivers of Reactions) Flashcards

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1
Q

Law of Conservation of Energy

A

Energy can not be created or destroyed, but is exchanged between a system and its surroundings.

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2
Q

Units of Energy

A

1 megajoule
1000 kilojoules
1 000 000 joules

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3
Q

Bond Energy

A

The amount of energy required to break 1 mol of a gaseous bond into separate atoms under standard conditions.

Bond energies of compounds are always positive, and given in kilojoules (kJ) or kilocalories (kcal).

Increase as atoms move closer, with bonds forming when the maximum possible bond strength is met. If atoms are pushed closer past this point, they begin to repel each other, and bond strength decreases.

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4
Q

Endothermic Reactions

A

Gain net energy, usually in the form of heat from the surrounding environment.

Absorbed energy is used to break bonds and typically grow cold.

Positive ∆H

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5
Q

Exothermic Reactions

A

Release net energy, usually in the form of heat into the surrounding environment.

Energy is released when more stable bonds form and are typically hot.

Negative ∆H

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6
Q

Changes of State

A

Substances remain at the same temperature while freezing, melting, or boiling until all molecules have changed state. Following this, the temperature then continues increasing or decreasing.

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7
Q

Sublimation

Changes of State

A

Endothermic

Solid → Gas

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8
Q

Melting

Changes of State

A

Endothermic

Solid → Liquid

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9
Q

Evaporation

Changes of State

A

Endothermic

Liquid → Gas

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10
Q

Deposition

Changes of State

A

Exothermic

Gas → Solid

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11
Q

Condensation

Changes of State

A

Exothermic

Gas → Liquid

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12
Q

Freezing

Changes of State

A

Exothermic

Liquid → Solid

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13
Q

∆H

A

Enthalpy change or heat of reaction.

  • Released in an exothermic reaction (-) or absorbed in an endothermic reaction (+) by a system under constant pressure
  • Typically in kilojoules
  • Proportional to mole ratios and corresponds with direction of equation
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14
Q

Calculating Enthalpy Theoretically

A

∆H = H(products) - H(reactants)

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15
Q

Calculating Enthalpy Experimentally

A

Calorimetry - the use of a calorimeter (any device that measures energy changes in chemical and physical changes) to determine enthalpy

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16
Q

Constant Pressure Calorimetry

A

‘Coffee cup calorimetry’

  • Made using an insulative polystyrene cup to prevent heat loss
  • The substance can be added into this vessel, as well as the reactants and a thermometer to measure the temperature change
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17
Q

Constant Volume Calorimetry

A

‘Bomb calorimetry’

  • Uses a stainless-steel container, oxygen, spark, stirrer and a flammable substance, which are the same components of a bomb
  • Highly insulative and contained
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18
Q

Calculating the Efficiency of a Calorimeter

A

Percentage form (the higher efficiency, the less heat lost)

(expected heat energy/actual energy supplied) x 100

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19
Q

Heat of Combustion (∆Hc)

A

The change in enthalpy when a fuel undergoes combustion.

It is always negative because all combustion reactions are exothermic.

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20
Q

Heat Capacity (c)

A

Energy required to raise 1 gram of a substance by 1°C.

Measured in J g⁻¹ K⁻¹ OR J g⁻¹ °C⁻¹

Determined by intermolecular forces and the heat required to break them.

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21
Q

Heat Capacity of Water

A

4.184

Relatively high due to hydrogen bonds between molecules.

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22
Q

Enthalpy of Dissolution (∆Hsoln)

A

Heat released or absorbed at a constant pressure when a solute is dissolved in a solvent, disassociating the positive and negative ions in a solid ionic compound.

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23
Q

Catalysts

A

Provides an alternative pathway with a lower activation energy for a reaction to occur without altering the surface area, concentration or energy of reactants.

They have the same initial and final mass when participating in a reaction, meaning they are not written in the chemical equation.

More efficient in increasing rate of reaction than temperature, such as in ammonium harvesting and catalytic converters.

24
Q

Homogenous Catalyst

A

Same physical state as the reactants and products.

25
Q

Heterogenous Catalyst

A

Different physical state from the reactants and products of a reaction, making them more easily separated from the products and reused.

26
Q

The Haber Process

A

Used to produce ammonia gas for explosives in WWI, but is now used for fertilisers, nylons and pharmaceuticals.

Hydrogen and nitrogen gas once needed extremely high temperatures of over 3000°C to bond, but the Haber process uses powdered iron as a catalyst. Despite less collision occuring than at high temperatures, more successful collions occur.

  1. N₂ and H₂ molecules are adsorbed by the surface particles of iron powder.
  2. Existing covalent bonds are broken, making the molecules monoatomic.
  3. N and H atoms bond to form ammonia according the the following reaction: N₂₍g₎ + 3H₂₍g₎ → 2NH₃₍g₎

Adsorb - to gather a substance on a surface in a condensed layer

27
Q

Latent Heat

A

The heat required to change the physical state of a specified amount of a substance where the temperature does not continue to rise.

Includes the enthalpy of fusion and enthalpy of vaporisation.

28
Q

Enthalpy of Fusion

Latent Heat

A

The energy at melting point required to melt a solid into its liquid form. When heat is added to a solid, molecules gain kinetic energy, allowing them to move faster and further apart.

Measured in J mol⁻¹ OR kJ mol⁻¹

Molar enthalpy of fusion is the energy required to fuse one mole of a solid to its liquid form.
Specific enthalpy of fusion is the energy required to fuse one gram of a solid to its liquid form.

The enthalpy of fusion for melting is always endothermic (+), while solidification, the reverse state change, is always exothermic (-).

29
Q

Enthalpy of Vaporisation

Latent Heat

A

The energy required to transform 1 mol of a substance at its boiling point from liquid to gas. Heat energy increases the kinetic energy of liquid particles, so that intermolecular forces are overcome and particles can move more randomly and freely as a gas.

Measured in J mol⁻¹ OR kJ mol⁻¹

Molar enthalpy of vaporisation is the energy required to vaporise one mole of a liquid to its gaseous form.
Specific enthalpy of vaporisation is the energy required to vaporise one gram of a liquid to its gaseous form.

The heat of vaporisation is always endothermic (+), just at the enthalpy of condensation, the reverse change of state, is always exothermic (-).

30
Q

Calculating Enthalpy Change using Bond Energies

A

The bond energy of a chemical reaction can be found by adding the positive total bond energy of the reactants (energy for breaking bonds) to the negative total bond energy of the products (energy for making bonds).

31
Q

Hess’ Law

A

The energy released or absorbed in a reaction is constant, no matter the steps taken for the reactants to become the products.

This means that:

  1. Chemical reactions can be reversible if appropriate conditions are available.
  2. Energy is always released or absorbed in a chemical reactions. The enthalpy change (∆H) of a reverse equation is the enthalpy change of a forward reaction multiplied by -1.
  3. Equations for chemical reactions can be added and subtracted algebraically.
32
Q

Calculating Enthalpy Change Using Hess’ Law

A
  1. Reverse given equations and enthalpy changes so that the reactants and products are on the same sides as in the final equation.
  2. Multiply mole ratios and enthalpy changes of given equations so that they correspond with the final mole ratios.
  3. Write in an algorithm, and cancel elements or compounds that are present on both sides of the equations. The equation remaining should be that of the final equation.
  4. Add enthalpy change values of given equations to find the enthalpy change of the final equation.
33
Q

Open System

A

Energy and matter can be freely be exchanged between system and surroundings.

34
Q

Closed System

A

Energy, but not matter, can be freely exchanged between system and surroundings.

35
Q

Isolated System

A

Neither energy or matter can be exchanged between system or surroundings.

This is referred to by the term, adiabatic.

36
Q

First Law of Thermodynamics

A

Law of Conservation of Energy - Energy cannot be created or destroyed.

∆E (energy of an internal system) = q + W (work is all forms of energy other than heat)

37
Q

Second Law of Thermodynamics

A

For a spontaneous process, the entropy of the universe must increase.

  • Energy is released into the environment, so nature tends towards a state of higher entropy (higher disorder)
38
Q

Third Law of Thermodynamics

A

A perfect crystal at zero Kelvin has zero entropy.

39
Q

Potential Energy

A

Ep ≈ bond energy/other intermolecular bonds

40
Q

Entropy (S)

A

A thermodynamic quantity that is the measure of disorder in a system.

Can not be measured directly, but instead by change (∆S) in J K⁻¹ mol⁻¹.

Entropy is what drives spontaneous chemical reactions as natural systems progress toward the highest probability of disorder.

  • ∆Suniverse = ∆Ssystem + ∆Ssurroundings > 0
41
Q

Temperature

Factors Afftecting Change in Entropy

A

Directly proportional

An increasing temperature will increase the kinetic energy and movement of particles, so more positions and arrangements are possible. A decreasing temperature will decrease entropy.

42
Q

Atomic Size

Factors Afftecting Change in Entropy

A

Directly proportional

Larger atoms have more subatomic particles, resulting in more possible arrangements. There is also more space available for particles to take up.

43
Q

Chemical Complexity

Factors Afftecting Change in Entropy

A

Directly proportional

More intramolecular bonds allow particles to move in more directions. Particles in linear bonds are only able to move left and right, while trigonally bonded particles can rotate around the particle in the centre, vibrate, and stretch.

44
Q

Number of Moles

Factors Afftecting Change in Entropy

A

Directly proportional

A higher number of moles results in a higher number of atoms and subatomic particles. This corresponds with a more possible arrangements and space for particles to take up.

45
Q

Dissolution

Factors Afftecting Change in Entropy

A

Proportional

As dissociated ions spread throughout a liquid, there are more possible arrangements than as a solid mass.

46
Q

Allotropes

Factors Afftecting Change in Entropy

A

Higher in some than in others.

Allotropes that allow more movement of particles, such as with delocalised electrons in their structures, have greater entropy.

47
Q

Entropy of Graphite

A

Graphite has a hexagonal lattice, made up of sheets of carbon atoms bonded in a trigonal planer. Each carbon atom corresponds with a single delocalised electron.

  • Conductive: delocalised electrons allow electrical current
  • Soft, as carbon sheets are able to slide over each other
  • Therefore, higher entropy than diamond as delocalised electrons and carbon sheets are able to move.
48
Q

Entropy of Diamond

A

Diamond has rigid crystal lattice structure, with each carbon atom covalently bonded to four others. There are no delocalised electrons.

  • Does nto conduct electricity
  • Hard, rigid structure
  • Therefore, lower entropy than graphite as it is more ordered with less movement of particles.
49
Q

Spontaneous Reactions

A

Do not require energy to be added to the system for a reaction or physical change to occur, as they take energy from the surroundings.

Decreases the entorpy of the surroundings but increases the entropy of the system, increasing the entropy of the universe overall.

For example:

  • Spontaneous reactions can react as soon as they are mixed, causing a high increase in entropy and possibly producing gases.
  • Some spontaneous reactions that can occur at room temperature will occur even faster if provided with a spark (combustion)
50
Q

Non-Spontaneous Reactions

A

Require a constant energy input to occur

A negative or positive enthalpy change (∆H) does not indicate whether a reaction is spontaneous or non-spontaneous.

For example:

  • Non-spontaneous reactions do not react unless continually heated as they do not produce sufficient energy themselves.
  • Non-spontaneous reactions do not react even at constant high temperatures.
51
Q

Gibbs Free Energy (∆G)

A

A quantity used to determine whether a process is spontaneous through a measure of the effects of temperature, entropy change and enthalpy change.

∆G = ∆H - T∆S

If negative, the reaction will be spontaneous and vice-versa. If Gibbs free energy is 0, the reaction is in equilibrium. It is measured in kJ mol⁻¹.

52
Q

Is a reaction spontaneous when ∆S is positive and ∆H is negative?

A

Spontaneous at all temperatures

53
Q

Is a reaction spontaneous when ∆S is positive and ∆H is positive?

A

Spontaneous at high temperatures

54
Q

Is a reaction spontaneous when ∆S is negative and ∆H is negative?

A

Spontaneous at low temperatures

55
Q

Is a reaction spontaneous when ∆S is negative and ∆H is positive?

A

Always non-spontaneous

03 Chemistry (6 decks)

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