Chemical Energetics (topic 5/15)

Question 1

Density of water

Answer 1

1g cm^-3

Question 2

Bond enthalpy

Answer 2

The energy change when one mole of covalent bonds, in a gaseous molecule, is broken under standard conditions.

Question 3

Average bond enthalpy

Answer 3

The energy change when one mole of covalent bonds, in a gaseous molecule, is broken under standard conditions, averaged over similar compounds.

Question 4

System

Answer 4

The name given to the region or place where the chemical reaction is happening (for example, a test tube, or a conical flask).

Question 5

Surroundings

Answer 5

Are considered to be the rest of the universe.

Question 6

In an open system

Answer 6

The system can exchange both mass and energy with the surroundings.

Question 7

In a closed system

Answer 7

The system can exhange only energy with the surroundings.

Question 8

Enthalpy (H)

Answer 8

Is a measure of the amount of heat energy inside of a substance. It is stored in the chemical bonds and intermolecular forces in the form of potential energy.

Question 9

When heat is added to the system

Answer 9

The enthalpy (or heat content) increases.

Question 10

When heat is lost from the system

Answer 10

The enthalpy (or heat content) decreases.

Question 11

Endothermic reactions

Answer 11

Are reactions where energy is absorbed from the surroundings, so the ΔH of the reaction is considered to be positive, as energy enters the system.

Question 12

Exothermic reactions

Answer 12

Are reactions where energy is released to the surroundings, resulting in a ΔH which is negative, due to energy going out of the system.

Question 13

Hess’s law

Answer 13

States that regardless of the route through which a chemical reaction proceeds, the enthalpy change will always be the same, provided the initial and final states of the system remain the same.

Question 14

The standard enthalpy of formation of a substance (ΔHθf)

Answer 14

Is the enthalpy change that occurs when one mole of the substance is formed from its elements in their standard states and under standard conditions.

The ΔHθf of a substance tells us how stable the substance is compared to its elements, and allows us to find the enthalpy change of all reactions related to the substances and its elements.

The ΔHθf of an element in its most stable form is 0, since there is no chemical change when an element changes to an element.

Question 15

First ionization energy

Answer 15

Refers to the minimum energy required to remove one mole of electrons from one mole of gaseous atoms (generally used for cations)
ex. Na(g) -> Na+ (g) + e- (g)

Question 16

First electron affinity

Answer 16

Refers to the enthalpy change when one mole of gaseous electrons is added to one mole of gaseous atoms (generally used for anions)
ex. Cl(g) + e- (g) -> Cl- (g)

Question 17

Lattice Enthalpy (ΔHølat)

Answer 17

Refers to the enthalpy change that occures when one mole of a solid, ionic compound is seperated into gaseous ions under standard conditions.

The value of lattice enthalpy can be deduced from the enthalpy of formation for an ionic compound.

As ionic compounds have a very strong electrostatic force of attraction, an immense amount of energy is required to break them apart, and hence lattice enthalpies are highly endothermic.

Question 18

Born-Haber cycle

Answer 18

This method essentially splits up the formation of a lattice into several steps, adds the enthalpy changes of each step to calculate the overall enthalpy, then returns the negative of that value to obtain the lattice enthalpy

Question 19

Steps for Born-Haber cycle

Answer 19

1. Formation of gaseous atoms from solid atoms of the metal (enthalpy change of atomization)
2. Formation of monoatomic particles of the non-metal (bond enthalpy)
3. Formation of gaseous metal cation from gaseous metal atom (first ionization energy)
4. Formation of gaseous non-metal anion from monoatomic particle (first electron affinity)
5. Formation of lattice from the metal and non-metal ion (enthalpy of formation)

The lattice enthalpy of any compund equals 1+2+3+4-5, where each number corresponds to the energy change obtained in the respective step.

When asked to write equations to show the lattice enthalpy, ensure the reactants are monoatomic and gaseous.

Question 20

Ionic Model

Answer 20

1. Theoretical lattice enthalpies can be calculated using the ionic model.
2. The ionic model assumes that the only interaction between ions in a lattice is due to the electrostatic force of attraction betweeen them.
3. Hence the energy needed to seperate the ions is said to depends on the product of the ionic charges, and sum of the ionic radii.
4. Lattice enthalpy decreases with increasing ionic radius, and increases with increasing ionic charge

Question 21

Solution Enthalpy (ΔHøsol)

Answer 21

• Refers to the enthalpy change when one mole of a solute is dissolved in a solvent to infinite dilution, under standars conditions.
• Can be obtained experimentally by measuring enthalpy changes if solutions with increasing volumes of water, until a limit is reached.
• Ionic compounds dissolve very readily in water as they are strongly attracted to water as a polar solvent.
• Ions seperated from the lattice due to dissolution are said to be hydrated. The strength of interaction between the ions and the polar water molcules is given by their hydration enthalpies.
• Enthalpy of hydration refers to the enthalpy change when one mole of a gaseous ion is dissolved to form an infinitely dilute solution.
  ex. X+(g) -> X+(aq), Y-(g) -> Y-(aq)
• Hydration enthalpies of individual ions are difficult to measure directly, hence the hydration enthalpy of the H+ ion is used (details for this process are not required)
• H+(g) -> H+(aq), ΔH = -1130 kJ mol^-1
• Hydration enthalpies for all ions are negative, as they are caused by attractive forces to polar molecules and are exothermic.
• Hydration enthalpies are approximately inversly proportional to ionic radii, yet directly proportional to electrostatic attraction (dependent on charge)

Question 22

Entropy (S)

Answer 22

Refers to the distribution of available energy among particles. It is a way of measuring the degree of disorder of a system.

The more ways energy can be distributed in a system, the more disordered the system is, and the greater the entropy.

The total entropy of the universe increases with time.

Question 23

Changes in entropy

Answer 23

• The solid state is the most ordered, and the gaseous state the least.
• Hence, entropy increases as we move from solid to liquid to gas, and decreases as we move from gas to liquid to solid.
• Entropy also increases when the number of particles increases - for example, doubling the number of particles doubles the entropy.
• The change in entropy caused by a change in the number of gaseous particles is the largest, i.e. changes in the number of gaseous particles affect entropy more than any other factor.
• Changes in state and the number of moles are the two key factors used while predicting changes in entropy.

ex.
1. C2H5OH(l) -> C2H5OH(g) - increase in entropy as ethanol goes from liquid to solid
2. H2(g) + F2(g) -> 2HF(g) - entropy change is close to zero as both state and number of moles remain the same on both sides (2 moles of gas on both sides).

Question 24

Absolute Entropy

Answer 24

• Standard values of entropy can be calculated for each substance at standard temperature and pressure. This is called absolute entropy (S°)
• A perfectly ordered solid at absolute zero (-273° K) has zero entropy. All other states have greater, positive entropies.
• Therefore, entropies become more and more positive as we move from solid to liquid to gas.
• Entropy cannot be negative.

(Absolute entropy values for a range of substances can be found in Section 12 of the IB Data Booklet).

Question 25

Total entropy change

Answer 25

• The total entropy change of a reaction takes into account both the entropy change of the surroundings and the entropy change of the system.
• Exothermic reactions result in an increase in total entropy as heat is dispersed into the surroundings. ΔS (surroundings) is proportional to - ΔH (system)
• The temperature of the surroundings plays a role in determining the change in entropy - when an exothermic reaction takes place, already hot surroundings will undergo little change in entropy as opposed to cold surroundings. ΔS (surroundings) is proportional to 1 / T. It follows that: ΔS (surroundings) = - ΔH (system) / T
• The units for entropy are therefore J K^-1 mol^-1