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Chemistry > 13: Energetics II > Flashcards

13: Energetics II Flashcards

1
Q

Define lattice energy and its sign

A

The Enthalpy change when 1 mole of an ionic compound is formed from its gaseous ions under standard conditions (exo)

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2
Q

Why is lattice energy exothermic

A

Ions are combined to form ionic solid lattice, and bonds are formed releasing energy

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3
Q

What does the large negative enthalpy change of lattice energy suggest

A
  • It suggests that the ionic compound is more stable than its gaseous ions
  • This is due to the strong electrostatic forces of attraction between the oppositely charged ions
  • There are no forces of attraction between the gaseous ions, therefore they are less stable that the ions in the lattice
  • The more exothermic the value, the stronger the ionic bonds within the lattice
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4
Q

What is the enthalpy change of atomisation and its sign

A

The enthalpy change when 1 mole of gaseous atoms is formed from its element under standard conditions (endo)

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5
Q

Why is the enthalpy of atomisation endothermic

A

Energy is required to break the bonds between atoms, and break the element into its gaseous atoms

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6
Q

What is electron affinity and its sign

A

The energy change when 1 mole of electrons is gained by 1 mole of gaseous atoms of an element to form one mole of gaseous ions, under standard conditions (exo and endo)

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7
Q

Why is the first electron affinity exothermic

A

Energy is released when electrons are attracted to the atoms

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8
Q

Why is the second electron affinity endothermic

A

Due to the large force of repulsion that must be overcome between a negatively charged ion and the negatively charged electron, which requires energy

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9
Q

What is the first electron affinity of chlorine

A
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10
Q

What is the second electron affinity of oxygen

A
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11
Q

What is a Born-Haber cycle and it’s main principles

A
  • Born-Haber cycles enable the calculations of unknown enthalpies
  • Arrows for endothermic reactions go up as the energy of the substance increases
  • Arrows for the exothermic reactions go down as the energy of the substance decreases
  • The cycle shows all the steps from the Enthalpy of formation of the elements in their standard states, to the lattice energy
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12
Q

Label 1-6 and describe why their arrow is pointing up/down

A
  1. Enthalpy of formation of sodium chloride (arrow pointing down as bonds made, exothermic)
  2. Atomisation energy of sodium (arrow pointing up as bonds broken, endothermic)
  3. Atomisation energy of chlorine (arrow pointing up as bonds broken, endothermic)
  4. First ionisation energy of sodium (arrow pointing up as energy required to break force of attraction between electron and sodium, endothermic)
  5. First electron affinity of chlorine (arrow pointing down as energy released when electron is attracted, exothermic)
  6. Lattice Enthalpy of sodium chloride (arrow pointing down as bonds made, exothermic)
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13
Q

How to calculate an unknown Enthalpy from a Born-Haber cycle

A

Add/subtract the values beginning after the unknown Enthalpy, and ending before the unknown Enthalpy

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14
Q

What is important to remember when dealing with molecules that have 2 moles of an element in Born-Haber cycles (e.g. MgCl2)

A

Sometimes the Enthalpy has to be doubled/halfed

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15
Q

What are theoretical lattice energies

A

Values that assume a perfect ionic model, where the ions within are 100% ionic (spherical with purely electrostatic attractions)

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16
Q

What are experimental lattice energies

A

The actual measured lattice energies of ionic compounds, that differ from theoretical lattice energies

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17
Q

Why do theoretical and experimental lattice energies differ

A
  • Experimental usually have higher values
  • Due to the covalent character of the ions
  • Perfect ionic bonding doesn’t occur, due to the different polarising powers of the cations
  • More polarising power of cations = More covalent character
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18
Q

What affects polarising power

A
  • Charge density
  • Large charge density = Large polarising power
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19
Q

How does ionic radius of the cation affect charge density, and therefore polarising power

A
  • Smaller ionic radius = Greater charge density
  • As valence electrons are further away from the nuclei, forces of attraction decreases
  • Greater charge density = Greater polarising power
  • Due to the greater concentration of positive charge of the smaller area
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20
Q

How does the charge of the cation affect charge density, and therefore polarising power

A
  • Greater charge = Greater charge density
  • Due to the ion being able to form stronger electrostatic forces of attraction with the anion
  • Greater charge density = Greater polarising power
  • Due to greater forces of attraction between the cation and anion
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21
Q

What affects the ease in which an anion is polarised

A
  • Ionic radius
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22
Q

How does the ionic radius of an anion affect the ease in which its polarised

A
  • Large ionic radius = Easily polarised
  • Due to the forces of attractions between the valance electrons and the nuclei being weaker at larger ionic radii, causing the electron cloud of the anion to be easily distorted
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23
Q

What is ionisation energy and its sign

A

The amount energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous ions, under standard conditions (endo)

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24
Q

Why is ionisation energy endothermic

A

Due to the energy required to remove an electron

25
Q

What is the Enthalpy change of solution and it’s sign

A

The Enthalpy change when 1 mole of an ionic substance dissolves in sufficient water to form an infinitely dilute solution (exo and endo)

26
Q

Why can the Enthalpy change of solution be exothermic and endothermic

A
  • It depends on the chemical energy of the substance being dissolved
  • NaCl is dissolved endothermically
  • NaOH is dissolved exothermically
27
Q

What is the Enthalpy of hydration and its sign

A

The Enthalpy change when 1 mole of a specified gaseous ion dissolves in sufficient water to form an infinitely dilute solution (exo)

28
Q

How does Enthalpy of hydration work

A
  • The positive and negative ions formed when ionic solid is dissolved in water form ion-dipole attractions with the water molecules
  • This is due to water being a polar molecule, where the oxygen atom is attracted to the positive ions and the hydrogen atoms are attracted to the negative ions
29
Q

How are Enthalpy of solution and Enthalpy of hydration related

A

Enthalpy of solution = reverse lattice Enthalpy + hydration Enthalpy
- Where hydration Enthalpy is the sum of the hydration enthalpies of each ion
- If there is a molecule with different moles (e.g. MgCl2) this must be accounted for

30
Q

What do 1-3 represent (Born-Haber cycle can also be drawn)

A
  1. Enthalpy of hydration of potassium chloride
  2. Lattice energy of potassium chloride
  3. Enthalpy of solution of potassium and chloride ions
31
Q

What are the factors affecting lattice Enthalpy

A
  • Charge of ions making up the crystalline lattice
  • Radius of ions making up the crystalline lattice
32
Q

How does the ionic radius of ions affect lattice energy

A
  • Large ionic radius = Less exothermic lattice energy
  • Due to the charge on the ions being more spread out when ions are larger
  • Ions are also further apart from each other
  • Therefore electrostatic forces of attraction between oppositely charged ions in the lattice are weaker
33
Q

How does ionic charge affect the lattice energy

A
  • Increased ionic charge = More exothermic lattice energy
  • Due to the ions having higher charge density
  • Therefore there are stronger electrostatic forces of attraction between the oppositely charged ions in the lattice
  • More energy is released on the formation of the lattice
34
Q

How does charge density affect the Enthalpy of hydration

A
  • Higher charge density = Higher hydration Enthalpy
  • Due to the ions attracting the water molecules more strongly
35
Q

Define entropy

A

The entropy of a given system is the measure of how disordered/chaotic it is. A system becomes more disordered and its entropy increases, making it more energetically stable, and the natural direction of change is increasing total entropy

36
Q

Define standard molar entropy

A

The entropy per mole of a substance under standard conditions of 1atm and a stated temperature

37
Q

Why does entropy change when state changes

A
  • Solid compounds are ordered and have a low entropy
  • In liquid state, particles are still close together but are arranged randomly and can move freely, so are disordered and have a higher entropy
  • Gas has the highest entropy, as the particles are fully disordered
38
Q

Why does entropy change when moles of state in reactants and products change

A
  • Gas molecules increase entropy, and solid compounds decrease entropy
  • When the moles of solid/liquid/gas change from the reactants to the products, the entropy changes
  • E.g. if there is one mole of solid reactant and it decomposes to one mole solid and one mole of gas, the entropy has increased, as there is the addition of a mole of gas
  • E.g. ammonium carbonate and ethanoic acid, endothermic process where the entropy increases as a more moles of gas and solution are produced
39
Q

Why does entropy change when a solid is dissolved

A
  • When solids (e.g. ammonium nitrate) are dissolved in water, the temperature of the solution decreases (endothermic process)
  • The entropy of the system increases as the particles go from fixed position to disordered, and energy is taken into the system
  • However in ionic solids the ionic bonds break, increasing energy
  • But the ions form bonds to the solvent molecules, which decreases entropy
40
Q

Why does entropy change when two solids are reacted

A
  • When certain solids (e.g. barium hydroxide and ammonium chloride) are mixed, the temperature drops (endothermic process)
  • The entropy depends on the physical state of the compounds made
41
Q

How can reactions be spontaneous if there is a decrease in entropy

A

Entropy of surroundings and system (total entropy) must be considered

42
Q

What is the total entropy equation

A
43
Q

What are the units of entropy

A

J K^-1 mol^-1

44
Q

What is the equation for entropy change of system

A
45
Q

What is the equation for entropy of surroundings

A
46
Q

What is the feasibility if a reaction determined by

A
  • Enthalpy change
  • Entropy change
  • Therefore the Gibbs free energy
47
Q

What is the Gibbs equation and the units used

A
48
Q

What result of the Gibbs equation shows a feasible reaction

A

Delta G has to be equal or less than 0, for a reaction to be feasible

49
Q

How is the Gibbs equation rearranged to find the temperature at which a non-spontaneous reaction became feasible

A
50
Q

Including positive/negative values for delta H, delta S, and delta G, explain whether the reaction is spontaneous/feasible or not

A
51
Q

How is the Gibbs equation changed to be expressed as a straight line

A
52
Q

What does this graph of free energy against temperature show

A
  • The x-intercept shows where the reaction ceases to be spontaneous
  • Above that temperature delta G is positive
  • The y-intercept shows the reaction is exothermic
  • Exothermic reactions are favoured by lower temperatures, which is why delta G becomes increasingly negative as temperature falls
53
Q

How is the equilibrium constant (Kc) and Gibbs Free Energy linked

A
54
Q

How to rearrange delta G = -RTlnK to find K

A
55
Q

If the calculated equilibrium constant is more than one, what does it indicate about delta G

A
  • Products are favoured
  • Delta G is less than 0
56
Q

If the calculated equilibrium constant is equal to one, what does it indicate about delta G

A
  • Reaction is equilibrium
  • Neither reactants nor products are favoured
  • Delta G = 0
57
Q

If the calculated equilibrium constant is less than one, what does it indicate about delta G

A
  • Reactants favoured
  • Delta G is positive
58
Q

What are the limitations of Gibbs Free Energy

A
  • Can only be used to predict the feasibility of reactions under standard conditions
  • Just because a reaction is feasible doesn’t mean it will occur at an observable rate
  • Doesn’t take into account the kinetics of the reaction (rate)
  • The Ea may be too large for the reaction to occur
  • Although some reactions are less than 0 and feasible, they may not be kinetically feasible (e.g. decomposition of hydrogen peroxide)

Chemistry (19 decks)

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