Energetics (AQA A Level Chemistry): Flashcards

A diazonium salt is a compound containing the –N≡N⁺ group bonded to an aryl ring, formed by the reaction of a primary arylamine with nitrous acid (HNO₂) at 0–5 °C.

ΔH is the energy difference between reactants and products — it shows whether the reaction is exothermic or endothermic.

E_a is the energy difference between the reactants and the transition state — it determines how readily the reaction proceeds.

E_a (forward) = E_a (reverse) + Δ*H*

For an endothermic forward reaction (positive ΔH), the forward activation energy is greater than the reverse activation energy by the value of ΔH.

True.

The transition state is the highest-energy point on the reaction profile. It is always higher in energy than both reactants and products, which is why molecules need a minimum activation energy to reach it.

A Friedel–Crafts alkylation is an electrophilic substitution reaction in which an alkyl group is introduced into a benzene ring using a halogenoalkane and an AlCl₃ catalyst to generate a carbocation electrophile.

True.

The standard enthalpy of formation of any element in its standard state is defined as zero, because no change occurs when an element is "formed" from itself.

In an exothermic reaction, heat energy is released to the surroundings, the temperature of the environment increases, and ΔH is negative.

A coupling reaction is a reaction between a diazonium ion and an aromatic compound (such as phenol or an arylamine) to form a coloured azo compound, –N=N–, under alkaline or mildly acidic conditions.

A pressure of 100 kPa, a temperature of 298 K (25 °C), and all substances in their standard physical states.

In an endothermic reaction, heat energy is absorbed from the surroundings, the temperature of the environment decreases, the enthalpy of the system increases, and ΔH is positive.

True.

ΔH_neut^θ is defined for the formation of exactly one mole of water from an acid-base reaction. It is always exothermic because forming the O–H bond in water releases energy.

A benzene ring is a cyclic structure of six carbon atoms with delocalised π electrons spread equally across all six carbons, giving all C–C bonds equal length and making benzene unusually stable.

False.

A temperature rise indicates that heat has been released to the surroundings, so the reaction is exothermic and ΔH is negative. A temperature fall indicates an endothermic reaction with a positive ΔH.

q = m × c × Δ*T*

where q is heat transferred in joules, m is mass in grams and ΔT is temperature change in °C. c is the specific heat capacity of the solution (4.18 J g^-1 °C^-1 for water).

A Friedel–Crafts acylation is an electrophilic substitution reaction in which an acyl group (RCO–) is introduced into a benzene ring using an acyl chloride and an AlCl₃ catalyst to generate the electrophile RCO⁺.

ΔH = −q / n

where q is the heat transferred (in kJ) and n is the number of moles of the limiting reactant. The negative sign ensures the sign convention is correct: exothermic gives negative ΔH.

Step 1: ΔT = 6.8 °C

Step 2: mass of solution = 100.0 g

Step 3: q = 100.0 × 4.18 × 6.8 = 2842.4 J

Step 4: moles of HCl = 0.0500 mol

True.

Aqueous solutions are largely water, so the specific heat capacity of pure water is used as a reasonable approximation for the solution in calorimetry calculations.

A second electron affinity is the enthalpy change when one mole of electrons is added to one mole of gaseous 1− ions to form one mole of gaseous 2− ions. It is always endothermic (positive).

False.

For formation data, arrows point upward from the elements to each compound. A useful mnemonic is "form up, burn down" — combustion data arrows point downward to combustion products.

ΔH_r = Σ ΔH_f (products) − Σ ΔH_f (reactants)

Multiply each ΔH_f value by the number of moles of that substance in the balanced equation.

The sign must be reversed when the direction of travel around the cycle is opposite to the direction of the arrow on the diagram.

A first electron affinity is the enthalpy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous 1− ions. It is usually exothermic (negative).

ΔH_r = ΔH₂ − ΔH₁

ΔH₂ = enthalpy change along the direct route (elements → products)

ΔH₁ = enthalpy change along the indirect route (elements → reactants)

False.

The ΔH_f^θ of an element in its standard state is zero by definition, so no arrow needs to be drawn for elements in a formation-data cycle — they are simply placed at the bottom of the cycle.

1. Write the formation equation at the top; place combustion products below.

2. Draw downward arrows from each substance to its combustion products.

3. Multiply each ΔH_c value by the moles of that substance.

4. Apply: ΔH_f = Σ ΔH_c (elements) − ΔH_c (compound)

False.

For combustion data, arrows point downward from each substance to the combustion products. This follows the "form up, burn down" rule.

Formation equation: 3C (s) + 3H₂ (g) + ½O₂ (g) → CH₃COCH₃ (l)

ΔH_f = Σ ΔH_c (elements) − ΔH_c (compound)

= (3 × −394) + (3 × −286) − (−1821)

= −1182 − 858 + 1821 = −219 kJ mol⁻¹

Average bond enthalpies are calculated using Hess's Law cycles that combine:

1. Enthalpy changes of atomisation and formation.

2. Enthalpy changes of atomisation and combustion.

Each enthalpy value must be multiplied by the number of moles of that substance in the balanced equation. For example, if 2 mol of NaHCO₃ (s) reacts, the ΔH_f value is multiplied by 2.

Step 1: Write the balanced equation at the top.

Step 2: Write the constituent elements below the equation.

Step 3: Draw upward pointing arrows to each compound.

Step 4: Apply Hess's Law: ΔH_r = ΔH_f (products) − ΔH_f (reactants).

False.

Hess's Law is particularly useful precisely because many enthalpy changes cannot be measured directly. It allows ΔH_f to be calculated indirectly from experimentally measurable ΔH_c values, or vice versa.

The enthalpy change when one mole of gaseous atoms is formed from an element in its standard state under standard conditions. It is always endothermic.

A standard entropy change is a measure of the disorder (randomness) of a system at a given temperature. Entropy increases when a solid becomes a liquid or gas, or when the number of moles of gas increases.

The mean bond enthalpy is defined as the enthalpy change to break 1 mol of bonds in the gaseous state, averaged over a range of compounds/molecules.

False.

Bond enthalpy calculations use mean (averaged) values, whereas Hess's law uses actual enthalpy values specific to the compounds involved — so the two methods give different results.

ΔH = Σ(bond enthalpies of bonds broken) − Σ(bond enthalpies of bonds formed)

Bonds broken are endothermic (+); bonds formed are exothermic (−).

Bond enthalpy calculations assume all reactants and products are in the gaseous phase. If a liquid is involved, the result is inaccurate because the enthalpy of vaporisation is not accounted for.

The bond enthalpy of the same bond type varies depending on the molecular environment (surrounding atoms). Mean values are averaged across many different compounds to give a reliable estimate.

False.

Bond breaking is always endothermic (energy is absorbed). Bond forming is always exothermic (energy is released).