UNIT 5 - Energetics And Thermochemistry

Question 1

State the heat of formation of pure element in standard state.

Answer 1

Equal to 0.

Question 2

Define entropy [S].

Answer 2

Measure of disorder in the system
but also distribution of available energy among the particles.

Question 3

State the second law of thermodynamics in terms of entropy.

Answer 3

The total entropy of a system either increases or remains constant in any spontaneous process it never decreases.

Question 3

State the factors entropy depends on.

Answer 3

• temperature
• pressure

Question 3

State the formula for lattice enthalpy.

Answer 3

∆H= K (constant dependent on geometry) *
n (magnitude of charges on the ions) * m (magnitude of charges of the ions) / Rn+Rm (ionic radii).

Question 3

Determine with reason the feasibility of a reaction depending on the temperature (entropy, Gibbs).

Answer 3

Feasible at high temperature - all reaction with ∆S system because the ∆G is almost equal -T∆S, making ∆H negligible.

Feasible at low temperature - all exothermic reactions because at low temp the T∆S is almost equal to 0 thus making the ∆G=∆H

Question 3

Define lattice enthalpy.

Answer 3

Enthalpy change that occurs when one mole of a solid ionic compound is separated into gaseous ions under standard conditions.

Question 3

Define energy.

Answer 3

Measure of the ability to do work.

Question 3

State when does exothermic reaction occur in terms of entropy.

Answer 3

If the change of entropy of the system can compensate for the negative entropy change of the surroundings (as heat flows from the surrounding to the system).

Question 3

State how lattice enthalpy changes.

Answer 3

Decreases with increasing ion radius.

Increases with increasing ion charge.

Question 3

Define the enthalpy of combustion.

Answer 3

Energy change upon complete combustion of one mole of a substance.

Question 3

Define the specific heat capacity.

Answer 3

The amount of heat needed to raise the temperature of a unit mass of pure substance by 1C or 1K.

Question 3

State the formula for heat [q].

Answer 3

q = m * ∆T * C

Question 3

Define the enthalpy of atomization.

Answer 3

Enthalpy change that occurs when one mole of gaseous atoms is formed from the elements in their STP.

Question 3

State the 1st law of thermodynamics.

Answer 3

Energy cannot be neither created nor destroyed only converted.

Question 3

State the formula for entropy change of the reaction.

Answer 3

∆S =∑∆S prod - ∑∆S reactants

Question 3

State the formula for enthalpy change of a reaction using heat of formation.

Answer 3

∆H = ∑∆Hf products - ∑∆Hf substrates

Question 3

State the most ordered and most disordered state.

Answer 3

most ordered - solid state
most disorder - gaseous state

Question 3

Define enthalpy.

Answer 3

A measure of the amount of heat energy contained in a substance.

Question 4

State the formulae for the Gibbs free energy change of the reaction.

Answer 4

∆G = ∑∆G products - ∑∆G substrates

Question 5

List the possible systems and their definitions.

Answer 5

• open system - can exchange energy and matter with the surroundings
• closed system - can only exchange energy but not matter
• isolated system - perfect one no exchange of anything (neither energy nor matter)

Question 6

Define hydration enthalpy.

Answer 6

Strength of interaction between the polar water molecules and the separated ions.

Question 7

Define a system in terms of thermodynamics.

Answer 7

The area of interest.

Question 8

State the Hess law.

Answer 8

The enthalpy change for any chemical reaction is independent of the route, providing the starting, final conditions and reactants, products are the same.

Question 9

State the formulae and define the Gibbs free energy.

Answer 9

Potential of a system to do work.

∆G = ∆H - T∆S

Question 10

Define an enthalpy change of solution.

Answer 10

Enthalpy change when one mole of a solute is dissolved in a solvent to infinite dilution under STP.

Question 11

Determine spontaneity of a reaction depending on the values of ∆H, ∆S, ∆G and T.

Answer 11

• ∆H>0, ∆S>0

∆G < 0, T high - spontaneous
∆G > 0, T low - non-spontaneous

• ∆H>0, ∆S<0

∆G>0 - non-spontaneous

• ∆H<0, ∆S<0

∆G < 0, T low - spontaneous
∆G > 0, T high - non-spontaneous

• ∆H<0, ∆S>0

∆G < 0 - spontaneous

Question 12

State the formulae for enthalpy change using the enthalpies of combustion.

Answer 12

∆H = ∑∆Hc substrates - ∑∆H products

Question 13

State the formulae for enthalpy change in a reaction using bond enthalpy.

Answer 13

∆H = ∑E bonds broken - ∑E bonds formed

Question 14

State the spontaneity of the reaction depending on the value of Gibbs free energy.

Answer 14

∆G<0 - spontaneous (exothermic usually)
∆G>0 - non-spontaneous (endothermic)
∆G=0 - equilibrium

Question 15

List the factors the increase in temperature when heated depends on.

Answer 15

• mass of object
• heat added
• nature of the substance

Question 16

Define the surroundings in terms of thermodynamics.

Answer 16

Everything else (except for the system) in the universe.

Question 17

List the standard conditions for enthalpy changes.

Answer 17

• pressure of 100kPa
• concentrations of 1 mol/dm3 for all
• all substances in their standard states (298K,10^5Pa)

Question 18

Define the enthalpy change of formation.

Answer 18

The energy change upon a formation of one mole of substance from elements in their standard states and conditions.

Question 19

Define bond enthalpy.

Answer 19

Energy required to break one mole of bonds in covalent, gaseous molecules under STP, is an average for every molecule.

Question 20

State the formulae for heat capacity [C].

Answer 20

C = ∆heat/∆T