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a-level chemistry > Chapter 22 > Flashcards

Chapter 22 Flashcards

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AP® Chemistry flashcards Periodic Table flashcards
1
Q

Ionic bonds

A

Electrostatic attraction between oppositely-charged ions in the lattice structure

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2
Q

Lattice Enthalpy

A

Measure of the strength of ionic bonding in a giant ionic lattice

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3
Q

Lattice Enthalpy DEFINITION

A

Enthalpy change that accompanies the formation of one mole of an ionic compound from its gaseous ions under standard conditions

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4
Q

Equation for KCl lattice Enthalpy?

A

K+ (g) + Cl- (g) -> KCl (s)
Gaseous ions -> solid ionic compounds
It is an EXOTHERMIC change and will always be negative as it involves bond formation (which releases energy)

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5
Q

Depict lattice Enthalpy on a cycle?

A

It is a downwards arrow from gaseous ions to the ionic lattice

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6
Q

What is a born-Haber cycle?

A

Cannot be measured directly and must be calculated indirectly

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7
Q

Born-Haber route for KCl Route 1?

A

Elements in standard states -> gaseous atoms (endothermic - formation of gaseous atoms)
Gaseous atoms -> gaseous ions (endothermic - formation of gaseous ions)
Gaseous ions -> ionic lattice (LATTICE ENTHALPY - EXOTHERMIC)

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8
Q

Born Haber cycle Route 2?

A

Changing from standard state directly to the ionic lattice - this is the Enthalpy change of FORMATION - this is Exothermic (just one Enthalpy change)

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9
Q

Standard Enthalpy change of formation?

A

Enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions with all reactants and products in their standard states

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10
Q

Standard Enthalpy change of atomisation? Give example for Cl2

A

Enthalpy change that takes place for the formation of one mole of gaseous atoms from the element in its standard state under standard conditions
1/2Cl2 (g) -> Cl(g)

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11
Q

Nature of Enthalpy change of atomisation?

A

Always endothermic as bonds are broken to form gaseous atoms ; when the element is a gas in its standard state, the standard Enthalpy change of atomisation = bond Enthalpy of bond being broken
Cl2(g) -> 2Cl(g) = +242kJ/mol
1/2 Cl2(f) -> Cl(g) = 121kJ/mol

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12
Q

First ionisation energy?

A

Enthalpy change required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions
Na (g) -> Na+ (g) + e-

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13
Q

Nature of ionisation energies?

A

Endothermic because energy required to overcome the attraction between a negative electron and the positive nucleus

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14
Q

Electron affinity

A

Opposite of ionisation energy - measures the energy to GAIN electrons - first electron affinity is the Enthalpy change that takes place when one electron is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions
Cl(g) + e- -> Cl- (g)

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15
Q

Nature of electron affinity

A

Exothermic because the electron being added is attracted in towards the nucleus

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16
Q

Ionisation energy/electron affinity of NaCl?

A

Ionisation energy of Na to Na+ (endothermic) - release an electron
Electron affinity of Cl to Cl- (Exothermic) - gain that electron

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17
Q

Second electron affinity?

A

With Oxygen - second electron affinities are endothermic… a second electron is being gained by a negative ion which repels the electron away so energy must be put in to force the negatively charged electron onto the negative ion
First electron affinity Exothermic

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18
Q

What happens when salt dissolves in water?

A

Water molecules are able to break up the giant ionic lattice structure and overcome the strong electrostatic attractions between oppositely charged ions - this is what happens when salt dissolves in water

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19
Q

Standard Enthalpy of solution

A

Enthalpy change that takes place when one mole of a solute dissolves in a solvent ; if the solvent is water the ions from the ionic lattice finish up surrounded with water molecules as aqueous ions

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20
Q

Equation for standard Enthalpy change of solution of NaCl

A

Na+Cl- (s) + aq -> Na+ (aq) + Cl- (aq)

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21
Q

How can Enthalpy change of solution be determined experimentally?

A

Weigh a sample of the ionic lattice and pour a set amount of water into the plastic cup in the beaker (measure temperature) ; quickly tip all KCl into the water and stir until fully dissolved and the temperature no longer changes

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22
Q

Maths in experimentally determining Enthalpy change of solution

A

q = mc(change in temperature)
Where mass = mass of water (g) + mass of KCl
c = 4.18 (SHC)
Change in temperature
Figure out moles of KCl and then divide kJ of energy by moles

23
Q

What happens when a solid ionic compound dissolves in water?

A

Ionic lattice breaks up forming separate gaseous ions - opposite energy change from lattice energy which forms the ionic lattice from gaseous ions
Water molecules are attracted to and surround the ions - separate gaseous ions interact with polar water molecules to form hydrated aqueous ions… this energy change is called the Enthalpy change of hydration

24
Q

Enthalpy change of hydration

A

Enthalpy change that accompanies the dissolving of gaseous ions in water to form one mole of aqueous ions
NaCl (s) -> Na+(g) + Cl-(g) is endothermic you are breaking up the lattice (part of this is Enthalpy change of hydration (Exothermic or endothermic)
Exact opposite (just sign change) is formation of lattice

25
Q

What are three levels in Enthalpy change hydration cycles?

A

Gaseous ions
Aqueous ions
Ionic lattice

26
Q

Gaseous ions to ionic lattice

A

Lattice Enthalpy

27
Q

Gaseous ions to aqueous ions

A

Enthalpy change of hydration

28
Q

Ionic lattice to aqueous ions

A

Enthalpy change of solution

29
Q

Ionic compounds

A

High melting and boiling points
Soluble in polar solvents
Conduct electricity when molten or in aqueous solution

30
Q

Factors affecting lattice Enthalpy?

A

Effect of ionic size and effect of ionic charge

31
Q

Effect of ionic size on lattice Enthalpy?

A

Ionic radius increases
Attraction between ions decreases
Lattice energy less negative
Melting point decreases

32
Q

Effect of ion charge?

A

Ionic charge increases (best ratio 2:2 like CaO for example)
Attraction between ions increases
Lattice energy becomes more negative
Melting point increases

33
Q

Opposing effects with oppositely charged ions?

A

Na+ Mg2+ and Al3+
Increasing charge means more attraction
Decreasing size (more protons to same number of electrons)
S2- and Cl-
Increasing charge gives more attraction
Increasing size gives less attraction

34
Q

Uses of metal oxides?

A

Protective coating for the insides of furnaces

35
Q

What other factors needed for melting points?

A

Picking of ions

36
Q

Hydration Enthalpy factors?

A

Ionic size - as you go down the group, ionic radius increases and attraction between ion and water molecules decreases so hydration energy is less negative
Ionic charge - as you go along the period, ionic charge increases so attraction with water molecules increases and hydration energy becomes more negative

37
Q

Predicting solubility

A

If sum of hydration enthalpies is larger than the magnitude of lattice Enthalpy then the overall Enthalpy change will be Exothermic and the compound should dissolve

38
Q

What else does solubility depend on?

A

Temperature and entropy

39
Q

What is entropy?

A

A measure of disorder (S)

40
Q

Greater the entropy?

A

Greater the dispersal of energy and greater the disorder - tendency for energy to spread out
Greater the energy is spread out per kelvin per mole

41
Q

General entropy rules?

A

Solids have smallest entropies
Gases have greatest entropies

42
Q

If a system changes to become more random?

A

Energy is spread out more and entropy change is positive
Vice versa

43
Q

Solid to liquid to gas

A

Entropy increases

44
Q

Melting and boiling

A

Increases randomness and thus increases disorder and increases entropy

45
Q

Standard entropy?

A

Entropy of one mole of a substance under standard conditions (100kPa and 298 K) - they are ALWAYS POSITIVE and have units J/K/mol

46
Q

Calculating entropy changes?

A

Sum of Entropy of products - sum of entropy of reactants

47
Q

Feasibility?

A

A reaction can happen if products have a lower overall energy than the reactants - spontaneous is same as energetically feasible

48
Q

Free energy

A

Overall change in energy during a chemical reaction made of
1) Enthalpy change (heat transfer between chemical system and surroundings)
2) Entropy change at the temperature of the reaction ; dispersal of energy within the chemical system itself

49
Q

Gibbs equation

A

Enthalpy change with sureoundings - Temperature in Kelvin*Entropy of system

50
Q

If gibbs<0

A

Reaction is feasible

51
Q

Convert Celsius to Kelvin

A

+273

52
Q

What to remember in Gibbs energy?

A

Entropy should be divided by 1000 to become kJ/K/mol
As temperature increases, entropy becomes more significant

53
Q

Limitations of predictions made for feasibility

A

Free energy change is useful for predicting feasibility but many reactions have a negative Gibbs energy and do not seem to take place
BECAUSE IT DOES NOT TAKE INTO ACCOUNT ACTIVATION ENERGY - VERY SLOW RATE (NO ACCOUNT OF KINETICS OR RATE OF REACTION)

a-level chemistry (28 decks)

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