Andrew Voyles

IB Chemistry Assessment Statements: Topics 5 & 15: Energetics

5.1.1: Exothermic : A reaction which produces heat. Endothermic : A reaction which absorbs heat. Standard Enthalpy of reaction : The change in internal energy (H) through a reaction is ΔH.

5.1.2: All combustion and neutralization reactions are exothermic processes.

5.1.3: If the reaction produces heat (increases the temperature of the surroundings) then it’s exothermic, and ΔH is negative. If it decreases the temp (i.e. absorbs heat) then it’s endothermic, and ΔH is positive. Also, the yield of an equilibrium reaction which is exothermic will be increased if it occurs at low temps, and so for endothermic reactions at high temperatures.

5.1.4: The most stable state is where all energy has been released. Therefore, when going to a more stable state, energy will be released, and when going to a less stable state, energy will be gained. On an enthalpy level diagram, higher positions will be less stable (with more internal energy) therefore, if the product is lower, heat is released (more stable products, less stable reactants, ΔH is negative) but if it is higher, heat is gained (less stable products, more stable reactants, ΔH is positive).

5.2.1: Change in internal energy = mass x specific heat capacity x change in temperature (q = m x C x ΔT)

150.0 g of water (4.184 Jg-1oC-1), heated from 36.0 oC to 98.0 oC:

q=(150.0 g)(4.184 Jg-1oC-1)(98.0 oC – 36.0 oC) = 38.9 kJ

5.2.2: When a reaction is carried out in water, the water will gain or lose heat from (or to) the reaction, usually with little escaping the water. Therefore, the change in energy, and so the ΔH value, can be calculated with q = m x c x ΔT where q is equal to ΔH, m is the mass of water present, and c = 4.18 kJ Kg-1 K-1. This ΔH value can then be calculated back to find the enthalpy change for each mol of reactants. The solution should be placed in a container as insulated as possible,

to keep as much heat as possible from escaping. The temperature should be measured continuously , and the value used in the equation is the maximum change in temp from the initial position.

5.2.4: The solution should be placed in a container as insulated as possible, to keep as much heat as possible from escaping. The temperature should be measured continuously , and the value used in the equation is the maximum change in temp from the initial position. For a reaction in aqueous solution, it is assumed that the reaction goes to completion. For a combustion reaction, it is assumed that all of the hydrocarbon reactant combusts and that it is in a gaseous state.

5.3.1: Hess’ Law states that the total enthalpy change between given reactants and products is the same regardless of any intermediate steps (or the reaction pathway). To calculate:

Reverse any reactions which are going the wrong way and invert the sign of their delta-H values.

Divide or multiply the reactions until the intermediate products will cancel out when the reactions are vertically added (always multiply/divide the delta-H value by the same number).

Vertically add them. Divide or multiply the resulting reaction to the correct

coefficients.

5.4.1: Bond enthalpy : The enthalpy change when one mol of bonds are broken homolitically in the gas phase. i.e. X-Y(g) -> X(g) + Y(g) : ΔH(dissociation). Molecules such as CH4 have multiple C-H bonds to be broken, and so the bond enthalpy for C-H is actually an average value. These values can be used to calculate unknown enthalpy changes in reactions where only a few bonds are being formed/broken.

5.4.2: If the sum total of the energy of all the bonds formed is greater than the sum total of the energy of all the bonds broken, then the reaction will be exothermic. If the sum total of the energy of all the bonds formed is less than the sum total of the energy of all the bonds broken, then the reaction will be endothermic.

15.1.1: Standard state : 101 kPa, 298 K (or 1 atm, 25 degrees celcuis). Standard enthalpy change of formation : The enthalpy change when 1 mol of a substance is made from its elements in their standard states. For example C(graphite) + 2H2(g) -> CH4(g). Molecules, like H2 are considered

to be 'standard state'. Fractions of mols (i.e. fractions in coefficients), may also be used if necessary as 1 mol must be produced). L. L.Standard enthalpy change of combustion: The enthalpy change when one mole of a substance completely reacts by combustion with oxygen under standard state.

15.1.2: If a reaction can be expressed in terms of enthalpy changes of formation, then add up all the ΔH values to get the ΔH for the reaction.

15.2.1: Lattice enthalpy : The enthalpy change when 1 mol of crystals (i.e. an ionic lattice) is formed from its component particles at an infinite distance apart. M+

(g) + X-(g) -> MX(s)

The value of lattice enthalpy is assumed to be positive for the separation of the lattice, and negative for the formation of the lattice.

15.2.2: Lattice enthalpy increases with higher ionic charge and with smaller ionic radius (due to increased attraction).

15.2.4: A significant difference between the theoretical and experimental lattice enthalpy values of ionic compounds indicates covalent character.

15.3.1: Factors which increase disorder (entropy) in a system: Mixing of particles. Change of state to greater distance between particles (solid ->

liquid or liquid -> gas). Increased particle movement (temperature). Increased number of particles (when more gas particles are

produced, this generally outweighs all other factors).

15.3.2: Predict the sign of ΔS (the change in entropy) for a reaction based on the above factors. ΔS is positive when entropy increases (more disorder) and negative when entropy decreases (less disorder).

15.3.3: The standard entropy change can be calculated by subtracting the absolute entropy of the reactants from that of the products.

15.4.1: When ΔG is negative, the reaction is spontaneous, when it's positive, the reaction is not.

15.4.2: ΔG = ΔH - Temperature(in kelvin) x ΔS Spontaneity depends on ΔH, ΔS and the temperature at which the reaction takes place (or doesn't as the case may be).

15.4.3: There are four possibilities: 1. ΔG is always negative when ΔH is negative and ΔS is positive. 2. ΔG is negative at high temperatures if ΔH is positive and ΔS is

positive (i.e. an endothermic reaction is spontaneous when T x ΔS is greater than ΔH).

3. ΔG is negative at lower temperatures if ΔH is negative and ΔS is negative (exothermic reactions are spontaneous if ΔH is bigger than T x ΔS).

4. ΔG is never negative if ΔH is positive and ΔS negative.