Define standard enthalpy change (Δ H Ꝋ ).

A standard enthalpy change is the enthalpy change measured under standard conditions: a pressure of 100 kPa , a temperature of 298 K (25 °C), and each substance in its standard physical state.

True or False? Standard conditions for enthalpy measurements are 298 K and 100 kPa.

True. Standard conditions are a temperature of 298 K (25 °C) and a pressure of 100 kPa , with each substance in its standard physical state .

Define standard enthalpy of formation (Δ H f Ꝋ ).

The standard enthalpy of formation is the enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions.

Standard conditions for thermodynamic measurements are a pressure of .......... and a temperature of .......... K.

Standard conditions for thermodynamic measurements are a pressure of 100 kPa and a temperature of 298 K.

True or False? The standard enthalpy of formation of any element in its standard state is zero.

True. By definition, no enthalpy change occurs when an element is formed from itself. So Δ H f Ꝋ for O 2 (g), Fe (s), and all other elements in standard states is 0 kJ mol −1 .

Define standard enthalpy of combustion (Δ H c Ꝋ ).

The standard enthalpy of combustion is the enthalpy change when one mole of a substance is completely burnt in excess oxygen under standard conditions. It is always exothermic , so Δ H c Ꝋ is always negative .

What does the symbol Ꝋ (plimsoll) indicate in an enthalpy expression?

The Ꝋ symbol indicates that the enthalpy change was measured under standard conditions : 100 kPa pressure, 298 K, and all substances in their standard physical states .

What is the standard enthalpy of neutralisation (Δ H neut Ꝋ )?

The standard enthalpy of neutralisation is the enthalpy change when one mole of water is formed from the reaction between an acid and an alkali under standard conditions.

True or False? The standard enthalpy of formation of CO 2 (g) is zero.

False. Δ H f Ꝋ = 0 only for elements in their standard states. CO 2 is a compound , so its Δ H f Ꝋ is non-zero (it is approximately −393.5 kJ mol −1 ).

Define enthalpy ( H ).

The enthalpy of a substance is the total chemical energy stored within it. In a reaction, Δ H represents the enthalpy change between reactants and products.

True or False? An exothermic reaction has a negative enthalpy change, Δ H .

True. In an exothermic reaction the products have less energy than the reactants , so energy is released to the surroundings and Δ H is negative .

What happens to the temperature of the surroundings in an exothermic reaction?

In an exothermic reaction, heat energy is released to the surroundings, so the temperature of the surroundings increases . The energy of the system decreases.

In an exothermic reaction, Δ H is .......... In an endothermic reaction, Δ H is ..........

In an exothermic reaction, Δ H is negative . In an endothermic reaction, Δ H is positive .

True or False? Activation energy is shown on enthalpy level diagrams.

False. Activation energy is shown on reaction profile diagrams , not on enthalpy level diagrams. Enthalpy level diagrams only show the relative energy levels of reactants and products .

Define exothermic reaction .

An exothermic reaction is one in which the products have less enthalpy than the reactants , so heat energy is released to the surroundings. Δ H is negative and the surroundings temperature increases .

In an endothermic reaction, which has greater enthalpy — the reactants or the products?

In an endothermic reaction, the products have greater enthalpy than the reactants . Heat energy is absorbed from the surroundings, so Δ H is positive and the surroundings temperature decreases .

In thermodynamics, the .......... refers to the reacting substances themselves, while the .......... refers to everything else (e.g. the flask and the room).

In thermodynamics, the system refers to the reacting substances themselves, while the surroundings refers to everything else (e.g. the flask and the room).

True or False? In an endothermic reaction, the enthalpy change Δ H is positive.

True. In an endothermic reaction, heat is absorbed from the surroundings, so the energy of the system increases . This corresponds to a positive Δ H .

What is calorimetry ?

Calorimetry is the experimental measurement of enthalpy changes in chemical reactions using a calorimeter (e.g. a polystyrene cup, vacuum flask or metal can).

True or False? The specific heat capacity of water is 4.18 J g −1 K −1 .

True. The specific heat capacity of water is 4.18 J g −1 K −1 , meaning 4.18 J of energy is needed to raise the temperature of 1 g of water by 1 K (or 1 °C).

Define specific heat capacity ( c ).

Specific heat capacity ( c ) is the energy required to raise the temperature of 1 g of that substance by 1 °C (or 1 K). The unit is J g −1 K −1 .

The energy transferred as heat in a calorimetry experiment is calculated using: q = .......... × .......... × Δ T

The energy transferred as heat in a calorimetry experiment is calculated using: q = m × c × Δ T where m is mass , c is specific heat capacity and Δ T is the temperature change .

True or False? A rise in temperature during a calorimetry experiment indicates a negative Δ H .

True. A temperature rise means heat has been released to the surroundings, indicating an exothermic reaction with a negative Δ H .

What does Δ T represent in the equation q = mc Δ T ?

Δ T represents the temperature change during the experiment: Δ T = T final − T initial . It is measured in °C or K , and the magnitude is the same in both units.

To calculate the enthalpy change per mole from a calorimetry experiment: Δ H = .......... ÷ .......... where n is the number of moles of the limiting reactant.

To calculate the enthalpy change per mole from a calorimetry experiment: Δ H = q ÷ n where q is the energy transferred (in kJ) and n is the number of moles of the limiting reactant.

Why is Δ H c Ꝋ always given as a negative value?

Δ H c Ꝋ is always negative because combustion is an exothermic process. The products have less enthalpy than the reactants, so energy is released to the surroundings .

Why can aqueous solutions of acids, alkalis and salts be treated as water when calculating q in a calorimetry experiment?

Aqueous solutions are assumed to be largely water , so the specific heat capacity ( c = 4.18 J g −1 K −1 ) and mass of water are used as approximations without significant loss of accuracy.

True or False? A rise in temperature during a calorimetry experiment indicates an endothermic reaction.

False. A temperature rise in the surroundings indicates an exothermic reaction, where heat energy is released. An endothermic reaction causes the temperature to fall .

Hess's Law states that the total enthalpy change of a chemical reaction is independent of the route taken, provided the initial and final conditions are the same.

True or False? Hess's Law is based on the Law of Conservation of Energy.

True. Hess's Law is a consequence of the Law of Conservation of Energy : energy cannot be created or destroyed, only converted from one form to another.

Why is Hess's Law useful for reactions that cannot be studied directly in a laboratory?

Hess's Law allows the enthalpy change of a reaction to be calculated indirectly from other known enthalpy changes. For example, the enthalpy of formation of propane cannot be measured directly because carbon and hydrogen do not react under standard conditions.

Using standard enthalpies of formation, the enthalpy change of a reaction is calculated as: Δ r H = Δ H f ( .......... ) − Δ H f ( .......... )

Using standard enthalpies of formation, the enthalpy change of a reaction is calculated as: Δ r H = Δ H f ( products ) − Δ H f ( reactants )

True or False? The enthalpy change for a reaction depends on the route taken.

False. According to Hess's Law , the enthalpy change of a reaction is independent of the route taken. Only the initial and final states determine the overall Δ H .

In a Hess cycle using enthalpies of formation, which direction do the arrows point relative to the elements?

The arrows point upward from the elements to the compounds , showing that compounds are being formed from their elements. The elements are always placed at the bottom of the cycle.

In a Hess cycle, if you travel against the direction of an arrow, you must .......... the sign of the enthalpy change for that step.

In a Hess cycle, if you travel against the direction of an arrow, you must reverse (negate) the sign of the enthalpy change for that step.

Write the general equation for Δ r H using standard enthalpies of formation .

Δ r H Ꝋ = ΣΔ H f Ꝋ ( products ) − ΣΔ H f Ꝋ ( reactants ) Each value must be multiplied by the stoichiometric coefficient from the balanced equation.

Why can't the standard enthalpy of formation of propane be determined directly by experiment?

Carbon (graphite) and hydrogen (H 2 (g)) do not react together under standard conditions, so it is impossible to measure the enthalpy change for: 3C (s) + 4H 2 (g) → C 3 H 8 (g) A Hess cycle using combustion enthalpies is used instead.

True or False? The enthalpy change for a direct route equals the sum of enthalpy changes for an indirect route between the same start and end points.

True. This is a direct statement of Hess's Law : the total Δ H is the same regardless of the pathway taken, because energy is conserved.

Define bond enthalpy (bond dissociation enthalpy).

Bond enthalpy is the energy required to break one mole of a particular bond in a gaseous molecule under standard conditions. It is always an endothermic process, so bond enthalpies are always positive .

True or False? Bond breaking is an endothermic process.

True. Breaking a bond requires energy input from the surroundings, so it is endothermic and has a positive enthalpy value. The reverse process, bond formation , releases energy and is exothermic.

Define average bond enthalpy .

The average bond enthalpy is the energy needed to break one mole of a particular type of bond in a gaseous molecule, averaged over a range of similar compounds to account for the effect of different molecular environments .

In a bond enthalpy calculation: Δ r H θ = enthalpy change for bonds .......... + enthalpy change for bonds ..........

In a bond enthalpy calculation: Δ r H θ = enthalpy change for bonds broken + enthalpy change for bonds formed Note: bonds broken values are positive (endothermic); bonds formed values are negative (exothermic).

True or False? Bond formation is an endothermic process.

False. Bond formation releases energy to the surroundings and is exothermic , with a negative enthalpy value. It is bond breaking that requires energy input and is endothermic.

Why are average bond enthalpies used rather than exact bond dissociation enthalpies?

The enthalpy required to break a bond depends on the molecular environment (the other atoms present). An average is taken across a range of similar compounds to give a value that can be used generally in calculations.

In a bond enthalpy calculation, what sign do values for bonds broken and bonds formed carry?

Values for bonds broken carry a positive sign (endothermic, energy absorbed). Values for bonds formed carry a negative sign (exothermic, energy released).

If more energy is released when bonds are .......... than is required to .......... bonds, the overall reaction is exothermic .

If more energy is released when bonds are formed than is required to break bonds, the overall reaction is exothermic .

What physical state must molecules be in for bond enthalpy values to be defined?

Molecules must be in the gaseous state. Bond enthalpies are defined for bonds broken in gaseous molecules; using other states would introduce additional enthalpy changes such as enthalpy of vaporisation.

True or False? If more energy is released when bonds are formed than is required to break bonds, the reaction is exothermic.

True. The overall Δ H is the sum of energy in (bond breaking, positive ) and energy out (bond forming, negative ). If the bond-forming term dominates, Δ H is negative and the reaction is exothermic .

A LevelChemistryEdexcelFlashcards1. Physical ChemistryEnergetics I

Energetics I (Edexcel A Level Chemistry): Flashcards

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Define standard enthalpy change (ΔH^Ꝋ).

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Define standard enthalpy change (ΔH^Ꝋ).
A standard enthalpy change is the enthalpy change measured under standard conditions: a pressure of 100 kPa, a temperature of 298 K (25 °C), and each substance in its standard physical state.
True or False?
Standard conditions for enthalpy measurements are 298 K and 100 kPa.
True.
Standard conditions are a temperature of 298 K (25 °C) and a pressure of 100 kPa, with each substance in its standard physical state.
Define standard enthalpy of formation (ΔH_f^Ꝋ).
The standard enthalpy of formation is the enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions.
Standard conditions for thermodynamic measurements are a pressure of .......... and a temperature of .......... K.
Standard conditions for thermodynamic measurements are a pressure of 100 kPa and a temperature of 298 K.
True or False?
The standard enthalpy of formation of any element in its standard state is zero.
True.
By definition, no enthalpy change occurs when an element is formed from itself. So ΔH_f^Ꝋ for O₂ (g), Fe (s), and all other elements in standard states is 0 kJ mol⁻¹.
Define standard enthalpy of combustion (ΔH_c^Ꝋ).
The standard enthalpy of combustion is the enthalpy change when one mole of a substance is completely burnt in excess oxygen under standard conditions. It is always exothermic, so ΔH_c^Ꝋ is always negative.
What does the symbol Ꝋ (plimsoll) indicate in an enthalpy expression?
The Ꝋ symbol indicates that the enthalpy change was measured under standard conditions: 100 kPa pressure, 298 K, and all substances in their standard physical states.
What is the standard enthalpy of neutralisation (ΔH_neut^Ꝋ)?
The standard enthalpy of neutralisation is the enthalpy change when one mole of water is formed from the reaction between an acid and an alkali under standard conditions.
True or False?
The standard enthalpy of formation of CO₂ (g) is zero.
False.
ΔH_f^Ꝋ = 0 only for elements in their standard states. CO₂ is a compound, so its ΔH_f^Ꝋ is non-zero (it is approximately −393.5 kJ mol⁻¹).
Define enthalpy (H).
The enthalpy of a substance is the total chemical energy stored within it. In a reaction, ΔH represents the enthalpy change between reactants and products.
True or False?
An exothermic reaction has a negative enthalpy change, ΔH.
True.
In an exothermic reaction the products have less energy than the reactants, so energy is released to the surroundings and ΔH is negative.
What happens to the temperature of the surroundings in an exothermic reaction?
In an exothermic reaction, heat energy is released to the surroundings, so the temperature of the surroundings increases. The energy of the system decreases.
In an exothermic reaction, ΔH is ..........
In an endothermic reaction, ΔH is ..........
In an exothermic reaction, ΔH is negative.
In an endothermic reaction, ΔH is positive.
True or False?
Activation energy is shown on enthalpy level diagrams.
False.
Activation energy is shown on reaction profile diagrams, not on enthalpy level diagrams. Enthalpy level diagrams only show the relative energy levels of reactants and products.
Define exothermic reaction.
An exothermic reaction is one in which the products have less enthalpy than the reactants, so heat energy is released to the surroundings. ΔH is negative and the surroundings temperature increases.
In an endothermic reaction, which has greater enthalpy — the reactants or the products?
In an endothermic reaction, the products have greater enthalpy than the reactants. Heat energy is absorbed from the surroundings, so ΔH is positive and the surroundings temperature decreases.
In thermodynamics, the .......... refers to the reacting substances themselves, while the .......... refers to everything else (e.g. the flask and the room).
In thermodynamics, the system refers to the reacting substances themselves, while the surroundings refers to everything else (e.g. the flask and the room).
True or False?
In an endothermic reaction, the enthalpy change ΔH is positive.
True.
In an endothermic reaction, heat is absorbed from the surroundings, so the energy of the system increases. This corresponds to a positive ΔH.
What is calorimetry?
Calorimetry is the experimental measurement of enthalpy changes in chemical reactions using a calorimeter (e.g. a polystyrene cup, vacuum flask or metal can).
True or False?
The specific heat capacity of water is 4.18 J g⁻¹ K⁻¹.
True.
The specific heat capacity of water is 4.18 J g⁻¹ K⁻¹, meaning 4.18 J of energy is needed to raise the temperature of 1 g of water by 1 K (or 1 °C).
Define specific heat capacity (c).
Specific heat capacity (c) is the energy required to raise the temperature of 1 g of that substance by 1 °C (or 1 K). The unit is J g⁻¹ K⁻¹.
The energy transferred as heat in a calorimetry experiment is calculated using:
q = .......... × .......... × ΔT
The energy transferred as heat in a calorimetry experiment is calculated using:
q = m × c × ΔT
where m is mass, c is specific heat capacity and ΔT is the temperature change.
True or False?
A rise in temperature during a calorimetry experiment indicates a negative ΔH.
True.
A temperature rise means heat has been released to the surroundings, indicating an exothermic reaction with a negative ΔH.
What does ΔT represent in the equation q = mcΔT?
ΔT represents the temperature change during the experiment: ΔT = T_final − T_initial. It is measured in °C or K, and the magnitude is the same in both units.
To calculate the enthalpy change per mole from a calorimetry experiment:
ΔH = .......... ÷ ..........
where n is the number of moles of the limiting reactant.
To calculate the enthalpy change per mole from a calorimetry experiment:
ΔH = q ÷ n
where q is the energy transferred (in kJ) and n is the number of moles of the limiting reactant.
Why is ΔH_c^Ꝋ always given as a negative value?
ΔH_c^Ꝋ is always negative because combustion is an exothermic process. The products have less enthalpy than the reactants, so energy is released to the surroundings.
Why can aqueous solutions of acids, alkalis and salts be treated as water when calculating q in a calorimetry experiment?
Aqueous solutions are assumed to be largely water, so the specific heat capacity (c = 4.18 J g⁻¹ K⁻¹) and mass of water are used as approximations without significant loss of accuracy.
True or False?
A rise in temperature during a calorimetry experiment indicates an endothermic reaction.
False.
A temperature rise in the surroundings indicates an exothermic reaction, where heat energy is released. An endothermic reaction causes the temperature to fall.
Define Hess's Law.
Hess's Law states that the total enthalpy change of a chemical reaction is independent of the route taken, provided the initial and final conditions are the same.
True or False?
Hess's Law is based on the Law of Conservation of Energy.
True.
Hess's Law is a consequence of the Law of Conservation of Energy: energy cannot be created or destroyed, only converted from one form to another.
Why is Hess's Law useful for reactions that cannot be studied directly in a laboratory?
Hess's Law allows the enthalpy change of a reaction to be calculated indirectly from other known enthalpy changes. For example, the enthalpy of formation of propane cannot be measured directly because carbon and hydrogen do not react under standard conditions.
Using standard enthalpies of formation, the enthalpy change of a reaction is calculated as:
Δ_rH = ΔH_f(..........) − ΔH_f(..........)
Using standard enthalpies of formation, the enthalpy change of a reaction is calculated as:
Δ_rH = ΔH_f(products) − ΔH_f(reactants)
True or False?
The enthalpy change for a reaction depends on the route taken.
False.
According to Hess's Law, the enthalpy change of a reaction is independent of the route taken. Only the initial and final states determine the overall ΔH.
In a Hess cycle using enthalpies of formation, which direction do the arrows point relative to the elements?
The arrows point upward from the elements to the compounds, showing that compounds are being formed from their elements. The elements are always placed at the bottom of the cycle.
In a Hess cycle, if you travel against the direction of an arrow, you must .......... the sign of the enthalpy change for that step.
In a Hess cycle, if you travel against the direction of an arrow, you must reverse (negate) the sign of the enthalpy change for that step.
Write the general equation for Δ_rH using standard enthalpies of formation.
Δ_rH^Ꝋ = ΣΔH_f^Ꝋ(products) − ΣΔH_f^Ꝋ(reactants)
Each value must be multiplied by the stoichiometric coefficient from the balanced equation.
Why can't the standard enthalpy of formation of propane be determined directly by experiment?
Carbon (graphite) and hydrogen (H₂ (g)) do not react together under standard conditions, so it is impossible to measure the enthalpy change for:
3C (s) + 4H₂ (g) → C₃H₈ (g)
A Hess cycle using combustion enthalpies is used instead.
True or False?
The enthalpy change for a direct route equals the sum of enthalpy changes for an indirect route between the same start and end points.
True.
This is a direct statement of Hess's Law: the total ΔH is the same regardless of the pathway taken, because energy is conserved.
Define bond enthalpy (bond dissociation enthalpy).
Bond enthalpy is the energy required to break one mole of a particular bond in a gaseous molecule under standard conditions. It is always an endothermic process, so bond enthalpies are always positive.
True or False?
Bond breaking is an endothermic process.
True.
Breaking a bond requires energy input from the surroundings, so it is endothermic and has a positive enthalpy value. The reverse process, bond formation, releases energy and is exothermic.
Define average bond enthalpy.
The average bond enthalpy is the energy needed to break one mole of a particular type of bond in a gaseous molecule, averaged over a range of similar compounds to account for the effect of different molecular environments.
In a bond enthalpy calculation:
Δ_rH^θ = enthalpy change for bonds .......... + enthalpy change for bonds ..........
In a bond enthalpy calculation:
Δ_rH^θ = enthalpy change for bonds broken + enthalpy change for bonds formed
Note: bonds broken values are positive (endothermic); bonds formed values are negative (exothermic).
True or False?
Bond formation is an endothermic process.
False.
Bond formation releases energy to the surroundings and is exothermic, with a negative enthalpy value. It is bond breaking that requires energy input and is endothermic.
Why are average bond enthalpies used rather than exact bond dissociation enthalpies?
The enthalpy required to break a bond depends on the molecular environment (the other atoms present). An average is taken across a range of similar compounds to give a value that can be used generally in calculations.
In a bond enthalpy calculation, what sign do values for bonds broken and bonds formed carry?
Values for bonds broken carry a positive sign (endothermic, energy absorbed). Values for bonds formed carry a negative sign (exothermic, energy released).
If more energy is released when bonds are .......... than is required to .......... bonds, the overall reaction is exothermic.
If more energy is released when bonds are formed than is required to break bonds, the overall reaction is exothermic.
What physical state must molecules be in for bond enthalpy values to be defined?
Molecules must be in the gaseous state. Bond enthalpies are defined for bonds broken in gaseous molecules; using other states would introduce additional enthalpy changes such as enthalpy of vaporisation.
True or False?
If more energy is released when bonds are formed than is required to break bonds, the reaction is exothermic.
True.
The overall ΔH is the sum of energy in (bond breaking, positive) and energy out (bond forming, negative). If the bond-forming term dominates, ΔH is negative and the reaction is exothermic.

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