A-Level Chemistry (AQA) > 3.1.4 Energetics > Flashcards

3.1.4 Energetics Flashcards

Calculate and interpret enthalpy changes using calorimetry, Hess’s law, and bond enthalpy data. (62 cards)

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1
Q

What term describes a reaction that releases heat to the surroundings?

A

Exothermic

Temperature of the surroundings increases.

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2
Q

What term describes a reaction that absorbs heat from the surroundings?

A

Endothermic

Temperature of the surroundings decreases.

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3
Q

Fill in the blank:

The heat energy change of a reaction at constant pressure is called the _______ change.

A

enthalpy

Represented by the symbol ΔH.

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4
Q

Under what condition is enthalpy change measured?

A

Constant pressure

This is typical for reactions in open laboratory systems.

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5
Q

True or False:

An exothermic reaction has a negative enthalpy change.

A

True

Heat energy is released, so ΔH is negative.

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6
Q

True or False:

An endothermic reaction has a negative enthalpy change.

A

False

ΔH is positive because energy is absorbed.

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7
Q

What term describes enthalpy changes measured under standard conditions?

A

Standard enthalpy change

Written with the symbol °.

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8
Q

Fill in the blank:

Standard conditions include a pressure of _______ kPa.

A

100

A specified temperature is also stated, often 298 K.

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9
Q

What does the symbol ΔH°₍₂₉₈₎ indicate about a reaction?

A

Standard enthalpy change at 298 K

Measured at 298 K and 100 kPa.

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10
Q

What is the standard enthalpy of combustion?

A

The enthalpy change when one mole of a substance burns completely in oxygen under standard conditions.

Symbol ΔcH°.

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11
Q

What does ΔcH° represent?

A

The standard enthalpy of combustion.

Enthalpy change when one mole of a substance burns completely in oxygen under standard conditions.

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12
Q

What is the standard enthalpy of formation?

A

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

Symbol ΔfH°.

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13
Q

What does ΔfH represent?

A

The standard enthalpy of formation.

Enthalpy change when one mole of a compound is formed from its elements in their standard states?

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14
Q

True or False:

The standard enthalpy of formation of an element in its standard state is zero.

A

True

This is defined as zero by convention.

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15
Q

Why must elements be in their standard states when defining standard enthalpy of formation?

A

To ensure a consistent reference state.

Ensures enthalpy data can be compared and used reliably.

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16
Q

What experimental technique is used to measure heat energy changes in chemical reactions?

A

Calorimetry

Often carried out using an insulated container and thermometer.

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17
Q

Fill in the blank:

The heat energy change in a reaction is represented by the symbol _______.

A

q

Measured in joules (J).

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18
Q

What equation is used to calculate the heat energy change during a reaction?

A

q = mcΔT

m = mass, c = specific heat capacity.

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19
Q

Fill in the blank:

In the equation q = mcΔT, the symbol m represents the _______.

A

mass

Usually the mass of the solution in grams.

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20
Q

In the equation q = mcΔT, what does c represent?

A

Specific heat capacity

Energy required to raise temperature of 1 g by 1 K.

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21
Q

In the equation q = mcΔT, what does ΔT represent?

A

Temperature change

Final temperature minus initial temperature.

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22
Q

True or False:

The unit of heat energy change (q) is joules.

A

True

Often converted to kJ for enthalpy calculations.

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23
Q

How is the molar enthalpy change obtained after calculating q in calorimetry?

A

Divide by moles

ΔH = −q ÷ number of moles reacting.

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24
Q

Fill in the blank:

In calorimetry calculations, mass must usually be expressed in _______.

A

grams

Especially when using specific heat capacity values.

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25
Why is the **value of q** divided by the number of moles of reactant?
To obtain molar enthalpy change. ## Footnote Gives ΔH in kJ mol⁻¹.
26
# True or False: If the temperature increases during a reaction, the reaction is **exothermic**.
True ## Footnote Heat is released to the surroundings.
27
What sign does **ΔH** have for an **exothermic reaction** measured by **calorimetry**?
Negative ## Footnote Energy is released by the system.
28
What sign does **ΔH** have for an **endothermic reaction** measured by **calorimetry**?
Positive ## Footnote Energy is absorbed by the system.
29
What type of **container** is often used in **simple calorimetry experiments** to reduce heat loss?
Insulated container ## Footnote Commonly a polystyrene cup.
30
# Fill in the blanks: When reporting **calculated values from calorimetry experiments**, results should match the _\_\_\_\_\_\_ _\_\_\_\_\_\_ of the least precise measurement.
significant figures ## Footnote Avoids implying unrealistic precision.
31
What **principle** states that the **total enthalpy change** of a reaction is independent of the route taken?
Hess’s law ## Footnote Enthalpy change depends only on the initial and final states.
32
# Fill in the blanks: **Hess’s law** states that the **total enthalpy change** depends only on the \_\_\_\_\_\_\_ and \_\_\_\_\_\_\_ states.
initial; final ## Footnote The pathway of the reaction does not affect ΔH.
33
# True or False: **Hess’s law** can be used to determine enthalpy changes that are difficult to measure directly.
True ## Footnote Indirect routes allow calculation of ΔH values.
34
What type of **enthalpy data** can be used with **Hess’s law** to calculate reaction enthalpy changes?
Enthalpies of combustion ## Footnote Often used when formation data are unavailable.
35
What other type of **standard enthalpy data** is commonly used in **Hess’s law calculations**?
Enthalpies of formation ## Footnote Standard values from data tables.
36
# Fill in the blanks: In **Hess’s law cycles**, equations may be \_\_\_\_\_\_\_ or \_\_\_\_\_\_\_ to obtain the desired reaction.
reversed; multiplied ## Footnote The enthalpy values must be adjusted accordingly.
37
What must be done to the **enthalpy change** when a **chemical equation is reversed** in a Hess’s law cycle?
Change sign ## Footnote Exothermic becomes endothermic and vice versa.
38
What must be done to the **enthalpy value** when an **equation is multiplied by a factor**?
Multiply enthalpy value ## Footnote ΔH scales with the equation coefficients.
39
# True or False: Hess’s law calculations are based on the **conservation of energy**.
True ## Footnote Energy cannot be created or destroyed.
40
What type of **diagram** is commonly used to **visualise Hess’s law calculations**?
Hess cycle ## Footnote Shows alternative reaction routes.
41
# Fill in the blank: The **enthalpy change calculated using Hess’s law** is usually expressed in _\_\_\_\_\_\_ per mole.
kJ mol⁻¹ ## Footnote Standard unit for molar enthalpy changes.
42
Why can **Hess’s law** be applied to **enthalpy changes**?
Enthalpy is a state function. ## Footnote It depends only on initial and final states.
43
# True or False: **Hess’s law** can be used to calculate **enthalpy changes for reactions such as thermal decomposition**.
True ## Footnote Useful when direct measurement is difficult.
44
What **property of enthalpy** allows the **enthalpy changes of multiple steps to be added together**?
state function ## Footnote The total change equals the sum of individual steps.
45
What is the **energy** required to **break one mole of a specified bond** in gaseous molecules called?
Mean bond enthalpy ## Footnote Measured in the gaseous state and averaged over many compounds.
46
Why are **bond enthalpies** described as **mean values**?
Average across compounds ## Footnote The bond energy varies slightly in different molecules.
47
# True or False: **Bond enthalpy values** apply to bonds in the **gaseous state**.
True ## Footnote Intermolecular forces in liquids and solids would affect values.
48
What **unit** is used to express **bond enthalpy values**?
kJ mol⁻¹ ## Footnote Energy required per mole of bonds broken.
49
# Fill in the blank: **Breaking bonds** requires _\_\_\_\_\_\_.
energy ## Footnote Bond breaking is always endothermic.
50
# Fill in the blank: **Forming bonds** releases _\_\_\_\_\_\_.
energy ## Footnote Bond formation is exothermic.
51
What **general equation** is used to **estimate enthalpy change** using bond enthalpies?
Bonds broken minus bonds formed ## Footnote ΔH ≈ Σ(bonds broken) − Σ(bonds formed).
52
# True or False: Bond enthalpy calculations provide **exact values** for reaction enthalpies.
False ## Footnote They only give approximate values.
53
Why do bond enthalpy calculations only give **approximate values of ΔH**?
Bond enthalpies are averages. ## Footnote Actual bond energies vary between molecules.
54
Why might values calculated using **bond enthalpies** differ from values obtained using **Hess’s law**?
Hess’s law uses exact data. ## Footnote Bond enthalpies are mean values.
55
# Fill in the blank: In **bond enthalpy calculations**, energy required to break bonds is always _\_\_\_\_\_\_ in value.
positive ## Footnote Energy must be supplied to break bonds.
56
What **sign** is associated with **energy released when bonds are formed**?
Negative ## Footnote Energy is released to the surroundings.
57
# True or False: Bond enthalpy calculations are usually applied to reactions in the **gaseous phase**.
True ## Footnote Bond enthalpy data are defined for gaseous molecules.
58
What does **x** represent in the diagram? ## Footnote Note the directions of the arrows.
Products
59
How do you calculate ΔH1 in this diagram? ## Footnote Note the directions of the arrows.
ΔH1 = ΔH2 - ΔH3
60
How do you calculate ΔH1 in this diagram? ## Footnote Note the directions of the arrows.
ΔH1 = ΔH2 + ΔH3
61
How do you calculate ΔH1 in this diagram? ## Footnote Note the directions of the arrows.
ΔH1 = ΔH3 - ΔH2
62
What does **X** represent in this diagram? ## Footnote Note the directions of the arrows.
Combustion products
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