A student burns a small piece of magnesium ribbon in a Bunsen flame. The magnesium burns brightly and the surrounding air becomes noticeably warmer. What type of reaction is this?
Burning is a combustion reaction. The surrounding air warms up, which means thermal energy has moved from the reacting magnesium and oxygen out into the surroundings (the air). A reaction that transfers thermal energy to the surroundings is exothermic, so the answer is A.
Question 2
CoreExothermic and endothermic reactions[1]
A student dissolves solid potassium nitrate in water in a test tube. The temperature of the solution falls from 20°C to 12°C. What type of process is this?
As potassium nitrate dissolves, the temperature drops from 20°C to 12°C. A temperature drop in the surroundings means the process has absorbed thermal energy from those surroundings (the water). Taking in thermal energy from the surroundings, causing a temperature decrease, is the definition of endothermic, so the answer is B.
Question 3
CoreExothermic and endothermic reactions[1]
When the two components of an epoxy adhesive are mixed, they react together and the mixture becomes noticeably warm as it hardens. What type of reaction is taking place?
When the adhesive hardens, chemical bonds form between the two components. The mixture becoming warm shows that thermal energy is being released into the surroundings during this process. A reaction that releases thermal energy to the surroundings, raising its temperature, is exothermic, so the answer is A.
Question 4
CoreExothermic and endothermic reactions[1]
The reaction pathway for the complete combustion of carbon can be described using two energy values, on the same relative energy scale: the reactants (carbon and oxygen) have an energy of 0 kJ/mol, and the products (carbon dioxide) have an energy of −394 kJ/mol. What is the value of ∆H for this reaction?
∆H is defined as the energy of the products minus the energy of the reactants. Here, ∆H = (−394) − (0) = −394 kJ/mol. The negative sign shows the products have less energy than the reactants, confirming this is an exothermic reaction (as expected for combustion), so the answer is C.
Question 5
SupplementExothermic and endothermic reactions[1]
Petrol (mainly octane) burns inside a car engine. What is the sign of ∆H for this combustion reaction?
Burning petrol releases thermal energy as heat (and does useful work in the engine), so it is exothermic. By definition, ∆H is negative for exothermic reactions, so the answer is D.
Question 6
SupplementExothermic and endothermic reactions[1]
Limestone (calcium carbonate) is heated strongly in an industrial lime kiln:
CaCO3(s) → CaO(s) + CO2(g)
The reaction only continues while the limestone is heated. What is the sign of ∆H for this reaction?
Because the limestone only decomposes while it is being heated, thermal energy must constantly be absorbed from the Bunsen burner (or kiln burners) for the reaction to continue. This makes it endothermic, and endothermic reactions have ∆H positive, so the answer is B.
Question 7
SupplementExothermic and endothermic reactions[1]
A catalytic converter contains platinum and rhodium, which speed up the conversion of carbon monoxide to carbon dioxide in a car's exhaust:
2CO(g) + O2(g) → 2CO2(g)
Which statement correctly defines the activation energy of this reaction?
Activation energy, Ea, is defined as the minimum energy that colliding particles must have in order to react. Statement B matches this definition exactly, so the answer is B. The other options describe different energy quantities: overall energy released (A), total bond energy of reactants (C), and enthalpy change (D).
Question 8
SupplementExothermic and endothermic reactions[1]
Hydrogen gas reacts with iodine vapour: H2(g) + I2(g) ⇌ 2HI(g). Without a catalyst, the reaction has an activation energy of 165 kJ/mol. In the presence of a platinum catalyst, the reaction follows an alternative pathway with an activation energy of 55 kJ/mol. What effect does the platinum catalyst have on the enthalpy change, ∆H, of the reaction?
A catalyst speeds up a reaction by lowering the activation energy needed for particles to react — here, from 165 kJ/mol down to 55 kJ/mol. However, a catalyst does not change the energy of the reactants or the products, so the overall enthalpy change, ∆H, stays exactly the same. The answer is A.
Question 9
SupplementExothermic and endothermic reactions[1]
Sulfur burns in oxygen to form sulfur dioxide:
S(s) + O2(g) → SO2(g)
Which statement correctly describes the energy changes during this reaction?
Breaking the O=O double bond in oxygen requires an input of energy, since energy must be supplied to pull bonded atoms apart — this is always an endothermic process. Forming the new S=O bonds in sulfur dioxide releases energy, since forming bonds is always exothermic. This matches option C.
Question 10
SupplementExothermic and endothermic reactions[1]
Chlorine gas in the upper atmosphere can be broken apart by ultraviolet light:
The Cl–Cl bond must be broken for one chlorine molecule to form two chlorine atoms. Breaking any bond requires energy to be absorbed, making this an endothermic process (the energy comes from the absorbed ultraviolet light). The answer is A.
Question 11
SupplementExothermic and endothermic reactions[1]
Methane reacts with chlorine in a free-radical substitution reaction:
CH4(g) + Cl2(g) → CH3Cl(g) + HCl(g)
This reaction is exothermic overall. What does this tell you about the bonds broken and the bonds formed?
For any reaction, ∆H = energy to break bonds in reactants − energy released forming bonds in products. If the reaction is exothermic overall, this value must be negative, meaning less energy was used breaking bonds than was released forming new ones. This matches option C.
Question 12
SupplementExothermic and endothermic reactions[1]
Hydrogen reacts with chlorine: H2(g) + Cl2(g) → 2HCl(g)
First find the energy needed to break the bonds in the reactants: one H–H bond (436) plus one Cl–Cl bond (242), giving 678 kJ/mol. Next find the energy released forming the bonds in the products: two H–Cl bonds are formed (since there are 2 mol of HCl), so 2 × 431 = 862 kJ/mol. Then ∆H = energy to break bonds − energy released forming bonds = 678 − 862 = −184 kJ/mol. Checking by working backwards: −184 = 678 − 862 ✓. The negative sign confirms this matches option C.
Question 13
SupplementExothermic and endothermic reactions[1]
A student adds yeast, which contains enzymes, to a solution of hydrogen peroxide, increasing the rate of its decomposition. What effect does the yeast have on the activation energy and on the enthalpy change, ∆H, of the reaction?
The enzymes in yeast act as a biological catalyst, providing an alternative pathway with a lower activation energy so that more particles can react successfully, which speeds up the decomposition. Since the reactants and products themselves are unchanged, the enthalpy change, ∆H, of the reaction is not affected by the catalyst. This matches option D.
Question 14
CoreExothermic and endothermic reactions[1]
A student adds excess zinc powder to 50 cm3 of copper(II) sulfate solution in an insulated polystyrene cup and stirs. The temperature rises from 22°C to 34°C over two minutes. What does this show about the reaction between zinc and copper(II) sulfate solution?
The temperature of the solution rises from 22°C to 34°C, showing that thermal energy is being transferred from the reacting mixture into the surroundings (the solution and the cup). A reaction that releases thermal energy to its surroundings, raising their temperature, is exothermic, so the answer is B.
Question 15
SupplementExothermic and endothermic reactions[1]
In the Haber process, the forward reaction N2(g) + 3H2(g) → 2NH3(g) has ∆H = −92 kJ/mol. What is ∆H for the reverse reaction (the decomposition of ammonia into nitrogen and hydrogen)?
The enthalpy change of the reverse reaction always has the same size but the opposite sign to the forward reaction, because the same amount of energy is involved, just transferred in the opposite direction. Since the forward reaction has ∆H = −92 kJ/mol, the reverse reaction (decomposition of ammonia) has ∆H = +92 kJ/mol. The answer is D.
Question 16
CoreExothermic and endothermic reactions[4]
A simple hand warmer contains iron powder in a sealed sachet. When the sachet is opened, air enters and reacts slowly with the iron powder:
4Fe(s) + 3O2(g) → 2Fe2O3(s)
The temperature of the sachet rises noticeably and stays warm for several hours.
(a)
State what is meant by the term exothermic reaction.
[2]
(b)
Using the hand warmer as an example, explain how you can tell that this reaction is exothermic.
a reaction that transfers thermal energy to the surroundings (1)
causing an increase in the temperature of the surroundings (1)
(b)[2]
the temperature of the sachet (the surroundings) rises / increases (1)
this shows that thermal energy is being released from the reacting iron and oxygen into the surroundings (1)
ALLOW converse phrasing, e.g. "energy is given out, so the temperature goes up"
Worked solution
An exothermic reaction is one that transfers thermal energy to the surroundings, causing the temperature of the surroundings to rise. In the hand warmer, iron reacts slowly with oxygen from the air to form iron(III) oxide. You can tell this reaction is exothermic because the sachet itself gets warm and stays warm for hours — this warmth is thermal energy that has been released by the chemical reaction into the surroundings (the sachet and the air around it, including the person holding it).
Question 17
SupplementExothermic and endothermic reactions[6]
An instant cold pack works by dissolving solid ammonium nitrate in water. A student places 5 g of ammonium nitrate and 20 cm3 of water in an insulated cup and stirs. The temperature of the water falls from 21°C to 9°C over one minute.
(a)
State what is meant by the term endothermic reaction.
[2]
(b)
State the sign of ∆H for the dissolving of ammonium nitrate in water, and explain your reasoning.
[2]
(c)
On a reaction pathway diagram for this process, state whether the energy of the products (the dissolved ions in solution) would be drawn above or below the energy of the reactants (the solid ammonium nitrate and water), giving a reason for your answer.
a reaction that takes in thermal energy from the surroundings (1)
causing a decrease in the temperature of the surroundings (1)
(b)[2]
positive / +ve (1)
because the process is endothermic — thermal energy is taken in from the surroundings, shown by the fall in temperature (1)
ALLOW ECF from (a) if the reasoning is internally consistent
(c)[2]
above / higher than the reactants (1)
because in an endothermic process the products end up with more energy than the reactants — energy has been absorbed into the products (1)
Worked solution
An endothermic reaction takes in thermal energy from the surroundings, cooling them down. Here, the temperature of the water falls from 21°C to 9°C, showing that dissolving ammonium nitrate absorbs thermal energy from the water. Since the process is endothermic, ∆H is positive. On a reaction pathway diagram, this means the products (the dissolved ions) are drawn above the reactants (the solid and water), because energy has flowed into the system to form them — the gap between the reactant and product energy levels represents the ∆H value, which is positive here.
Question 18
SupplementExothermic and endothermic reactions[6]
Propane burns completely in a camping stove. The reaction pathway for this combustion can be described using the following energy values, all on the same relative energy scale:
energy of the reactants (propane + oxygen) = 100 kJ
energy of the highest point on the pathway = 160 kJ
energy of the products (carbon dioxide + water) = 20 kJ
(a)
Calculate the activation energy, Ea, for the forward reaction.
[1]
(b)
Calculate the enthalpy change, ∆H, for the combustion of propane.
[2]
(c)
State whether the reaction is exothermic or endothermic, and explain how this can be seen from the position of the products relative to the reactants on the reaction pathway.
[2]
(d)
Calculate the activation energy for the reverse reaction (the decomposition of carbon dioxide and water back into propane and oxygen).
because the products are at a lower energy than the reactants, so energy has been released to the surroundings (1)
ECF from (b)
(d)[1]
Ea(reverse) = 160 − 20 = 140 kJ
Worked solution
The activation energy is the difference between the energy of the reactants and the highest point on the pathway: 160 − 100 = 60 kJ for the forward reaction. The enthalpy change is the difference between the energy of the products and the energy of the reactants: 20 − 100 = −80 kJ. Because this is negative, the products end up with less energy than the reactants, so the reaction is exothermic — the "extra" 80 kJ has been released to the surroundings. For the reverse reaction, the activation energy is measured from the products up to the same highest point: 160 − 20 = 140 kJ. Notice the reverse activation energy (140 kJ) is larger than the forward one (60 kJ) by exactly 80 kJ, which is the size of ∆H — this is always true for a reversible reaction.
Question 19
CoreExothermic and endothermic reactions[5]
Living cells constantly carry out two important reactions:
respiration: C6H12O6 + 6O2 → 6CO2 + 6H2O, which releases energy for the cell to use
photosynthesis: 6CO2 + 6H2O → C6H12O6 + 6O2, which requires energy absorbed from sunlight
(a)
State, with a reason, whether respiration is exothermic or endothermic.
[2]
(b)
State, with a reason, whether photosynthesis is exothermic or endothermic.
[2]
(c)
State one way in which the two reactions above are related to each other.
because energy (thermal energy) is released, transferred to the surroundings, during the reaction (1)
(b)[2]
endothermic (1)
because energy (thermal energy) must be absorbed, taken in from sunlight, for the reaction to occur (1)
(c)[1]
they are the reverse of each other / the reactants and products of one are the products and reactants of the other
ALLOW respiration releases the energy that photosynthesis originally stored in glucose
Worked solution
Respiration breaks down glucose using oxygen and releases energy that the cell can use for its life processes, so it is exothermic. Photosynthesis does the opposite: it uses energy absorbed from sunlight to build glucose from carbon dioxide and water, so it is endothermic. The two equations are exact reverses of one another, and the energy released by respiration is ultimately the same energy that was stored in the glucose during photosynthesis — energy is being passed from sunlight, into glucose, and then released again for the cell to use.
Question 20
SupplementExothermic and endothermic reactions[6]
When solid copper(II) carbonate is heated strongly in a test tube, it decomposes:
CuCO3(s) → CuO(s) + CO2(g)
The green solid turns black, and heating must continue for the reaction to carry on.
(a)
State, with a reason, whether this reaction is exothermic or endothermic.
[2]
(b)
Define the term activation energy.
[2]
(c)
Suggest and explain the effect that adding a suitable catalyst would have on (i) the activation energy and (ii) the enthalpy change, ∆H, of this reaction.
because thermal energy must continually be taken in (from the Bunsen burner) for the reaction to continue — shown by the need for constant heating (1)
(b)[2]
the minimum energy (1)
that colliding particles must have in order to react (1)
(c)[2]
(i) the activation energy would decrease / be lowered, because the catalyst provides an alternative reaction pathway (1)
(ii) the enthalpy change, ∆H, would stay the same / be unaffected (1)
Worked solution
Copper(II) carbonate only breaks down while it is being heated, showing that thermal energy must be continuously supplied for the reaction to keep going — this makes it endothermic. Activation energy is the minimum energy that colliding particles need in order to react. A catalyst works by offering an alternative reaction pathway with a lower activation energy, so more particles can react successfully at a given temperature. However, adding a catalyst never changes the energy of the reactants or products themselves, so it has no effect on the overall enthalpy change, ∆H, of the reaction.
Question 21
SupplementExothermic and endothermic reactions[7]
In margarine manufacture, hydrogen gas is added across the C=C double bond in ethene:
C2H4(g) + H2(g) → C2H6(g)
Use the bond energies in the table to answer the questions below.
Bond
Bond energy (kJ/mol)
C=C
612
C–C
347
C–H
413
H–H
436
(a)
State the bonds that are broken in this reaction, and calculate the total energy needed to break them.
[2]
(b)
State the bonds that are formed in this reaction, and calculate the total energy released when they form.
[2]
(c)
Use your answers to (a) and (b) to calculate ∆H for this reaction, and state whether it is exothermic or endothermic.
ALLOW counting every bond in the reactants, (4 × 413) + 612 + 436 = 2700 kJ (2)
(b)[2]
one C–C bond and two (new) C–H bonds are formed (1)
347 + (2 × 413) = 347 + 826 = 1173 kJ (1)
ALLOW counting every bond in the product, (6 × 413) + 347 = 2825 kJ (2), only if every reactant bond was also counted in (a)
(c)[3]
∆H = energy to break bonds − energy released forming bonds (1)
∆H = 1048 − 1173 = −125 kJ/mol (1)
exothermic (1)
ECF from (a) and (b)
ALLOW 2700 − 2825 = −125 kJ/mol from the count-every-bond method
REJECT mixing the two methods, e.g. 1048 − 2825 = −1777 kJ/mol; award the method mark only
Worked solution
In this reaction, the C=C double bond in ethene and the H–H bond in hydrogen are broken, while a new C–C single bond and two new C–H bonds are formed (the four C–H bonds already present in ethene are unchanged, so they are not counted). Breaking bonds needs 612 + 436 = 1048 kJ. Forming the new bonds releases 347 + (2 × 413) = 1173 kJ. So ∆H = 1048 − 1173 = −125 kJ/mol. Since more energy is released forming bonds than is needed to break them, the reaction is exothermic, which fits with hydrogenation reactions being used industrially without needing to add much extra energy.
Question 22
SupplementExothermic and endothermic reactions[6]
A hydrogen fuel cell combines hydrogen and oxygen gas to produce water vapour:
2H2(g) + O2(g) → 2H2O(g)
Use the bond energies in the table to answer the questions below.
Bond
Bond energy (kJ/mol)
H–H
436
O=O
496
O–H
463
(a)
Calculate the total energy needed to break all the bonds in the reactants.
[2]
(b)
Calculate the total energy released when all the bonds in the products form.
[2]
(c)
Calculate ∆H for this reaction, and state whether it is exothermic or endothermic.
2 mol of H2O contain 4 O–H bonds in total; 4 × 463 (1)
= 1852 kJ (1)
(c)[2]
∆H = 1368 − 1852 = −484 kJ/mol (1)
exothermic (1)
ECF from (a) and (b)
Worked solution
The reactants contain two H–H bonds and one O=O bond, so the energy needed to break them all is (2 × 436) + 496 = 1368 kJ. The products are 2 mol of H2O, and each water molecule has two O–H bonds, so 4 O–H bonds form in total, releasing 4 × 463 = 1852 kJ. The enthalpy change is ∆H = 1368 − 1852 = −484 kJ/mol. This is negative, so the reaction is exothermic, which is why hydrogen fuel cells can be used to generate useful energy.
Question 23
SupplementExothermic and endothermic reactions[6]
Hydrogen peroxide decomposes slowly at room temperature:
2H2O2(aq) → 2H2O(l) + O2(g)
Without a catalyst, this reaction has an activation energy of 75 kJ/mol. Adding manganese(IV) oxide as a catalyst provides an alternative pathway with an activation energy of 25 kJ/mol. In both cases, ∆H = −98 kJ/mol.
(a)
Explain what the value 75 kJ/mol represents for this reaction.
[2]
(b)
Explain, in terms of activation energy, why the manganese(IV) oxide catalyst increases the rate of decomposition.
[2]
(c)
State the value of ∆H when the manganese(IV) oxide catalyst is used, and explain why it is not affected by the catalyst.
the minimum energy that colliding particles (hydrogen peroxide molecules) must have in order to react (1)
(b)[2]
the catalyst provides an alternative reaction pathway with a lower activation energy (1)
so a greater proportion of particles have enough energy to react when they collide, which increases the rate of reaction (1)
(c)[2]
−98 kJ/mol (1)
because a catalyst only provides an alternative pathway (changes the activation energy) — it does not change the energy of the reactants or products, so ∆H is unaffected (1)
Worked solution
The 75 kJ/mol figure is the activation energy without a catalyst — the minimum energy that colliding hydrogen peroxide molecules need in order to react. Manganese(IV) oxide works as a catalyst by providing a different reaction pathway that needs only 25 kJ/mol of activation energy. Since more particles will have this lower amount of energy at a given temperature, more successful collisions happen per second, so the reaction goes faster. However, the catalyst does not alter the reactants or products, so the enthalpy change stays at −98 kJ/mol whether or not the catalyst is used.
Question 24
CoreExothermic and endothermic reactions[6]
A student adds 25 cm3 of dilute hydrochloric acid to 25 cm3 of dilute sodium hydroxide solution in an insulated cup, stirring and recording the temperature every 30 seconds.
Time (s)
0
30
60
90
120
Temperature (°C)
19.5
24.0
27.5
27.5
27.0
(a)
Calculate the increase in temperature between 0 s and its highest recorded value.
[1]
(b)
State what is meant by the term exothermic reaction.
[2]
(c)
Explain, using the data in the table, why the neutralisation reaction between hydrochloric acid and sodium hydroxide is described as exothermic.
a reaction that transfers thermal energy to the surroundings (1)
causing an increase in the temperature of the surroundings (1)
(c)[3]
the temperature of the mixture (the surroundings) increases, from 19.5°C to a maximum of 27.5°C (1)
this shows that thermal energy is being released by the reaction (1)
into the surroundings (the solution/cup), which matches the definition of an exothermic reaction (1)
ALLOW ECF from (a)
Worked solution
The highest recorded temperature is 27.5°C, so the temperature rise is 27.5 − 19.5 = 8.0°C. An exothermic reaction is one that transfers thermal energy to the surroundings, causing their temperature to rise. Because the temperature of the acid–alkali mixture rises as the two solutions react, thermal energy must be flowing out of the reacting particles and into the surroundings (the solution itself, which is being measured). This matches the definition of an exothermic reaction, confirming that neutralisation is exothermic.
Question 25
CoreExothermic and endothermic reactions[5]
Two products sold for sports injuries work in opposite ways:
a self-heating heat pack, which becomes warm when iron powder inside reacts with air
an instant cold spray, which becomes cold when a chemical inside dissolves in water
(a)
State which product contains a reaction that is exothermic, and explain your reasoning.
[2]
(b)
State which product contains a process that is endothermic, and explain your reasoning.
because thermal energy is transferred to the surroundings, shown by the temperature rise / it becoming warm (1)
(b)[3]
the cold spray (the chemical dissolving in water) (1)
because thermal energy is taken in from the surroundings (1)
shown by the temperature falling / it becoming cold (1)
Worked solution
The heat pack becomes warm because the reaction between iron and oxygen releases thermal energy into its surroundings — this release of energy, shown by the rise in temperature, is what makes a reaction exothermic. The cold spray becomes cold because the chemical dissolving inside it absorbs thermal energy from its surroundings (including your skin) rather than releasing it — this absorption of energy, shown by the fall in temperature, is what makes a process endothermic.
Question 26
SupplementExothermic and endothermic reactions[7]
In the Haber process, nitrogen and hydrogen react together:
N2(g) + 3H2(g) → 2NH3(g)
Use the bond energies in the table to answer the questions below.
Bond
Bond energy (kJ/mol)
N≡N
944
H–H
436
N–H
391
(a)
Calculate the total energy needed to break all the bonds in the reactants.
[2]
(b)
Calculate the total energy released when all the bonds in the products form.
[2]
(c)
Calculate ∆H for this reaction, and state whether it is exothermic or endothermic.
[2]
(d)
The reaction above is reversible. State whether the reverse reaction (the decomposition of ammonia into nitrogen and hydrogen) is exothermic or endothermic, and give the value of its enthalpy change.
ALLOW ECF from (c), with the sign reversed and the same size
Worked solution
Breaking the reactant bonds needs 944 (N≡N) + (3 × 436) (H–H) = 2252 kJ. Each ammonia molecule has 3 N–H bonds, so 2 mol of NH3 contain 6 N–H bonds, releasing 6 × 391 = 2346 kJ when they form. So ∆H = 2252 − 2346 = −94 kJ/mol, showing the forward reaction is exothermic. Because the reaction is reversible, the reverse reaction (ammonia decomposing back into nitrogen and hydrogen) must have exactly the opposite enthalpy change: +94 kJ/mol, making it endothermic.
Question 27
SupplementExothermic and endothermic reactions[6]
Sulfur dioxide reacts with oxygen in a reversible reaction:
2SO2(g) + O2(g) ⇌ 2SO3(g)
For the forward reaction, ∆H = −196 kJ/mol and the activation energy, Ea, is 250 kJ/mol.
(a)
State the value of ∆H for the reverse reaction (the decomposition of sulfur trioxide), giving a reason for your answer.
[2]
(b)
Calculate the activation energy for the reverse reaction.
[2]
(c)
State which direction of the reaction, forward or reverse, has the higher activation energy, and explain why, in terms of the energies of the reactants and products.
because the products (SO3) are at a lower energy than the reactants (the forward reaction is exothermic), so the reverse reaction must climb from this lower level to the same peak, i.e. Ea(reverse) = Ea(forward) + 196 (1)
ALLOW the reverse reaction is endothermic, so its activation energy includes the energy needed to raise the products up to the reactant level
Worked solution
Since the forward reaction is exothermic with ∆H = −196 kJ/mol, the reverse reaction must be endothermic with ∆H = +196 kJ/mol — the same size, opposite sign. On a reaction pathway diagram, the activation energy for the reverse reaction is measured from the (lower) product energy level up to the same peak that the forward reaction is measured from, so Ea(reverse) = Ea(forward) + 196 = 250 + 196 = 446 kJ/mol. Because the products are much lower in energy than the reactants (a large exothermic ∆H), the reverse reaction has to climb a much bigger energy hill, which is why its activation energy is so much higher than the forward reaction's.
Question 28
CoreExothermic and endothermic reactions[6]
A student dissolves separate 5 g samples of two different solids in 20 cm3 of water each, in insulated cups, and records the temperature change.
Solid
Starting temperature (°C)
Final temperature (°C)
Solid sodium hydroxide
20.0
32.5
Solid ammonium chloride
20.0
14.5
(a)
Calculate the temperature change for each solid.
[2]
(b)
State, with a reason, whether dissolving sodium hydroxide is exothermic or endothermic.
[2]
(c)
State, with a reason, whether dissolving ammonium chloride is exothermic or endothermic.
ALLOW a fall / decrease of 5.5°C without a minus sign
(b)[2]
exothermic (1)
because the temperature of the surroundings (the water) increases, showing thermal energy is being released to the surroundings (1)
(c)[2]
endothermic (1)
because the temperature of the surroundings (the water) decreases, showing thermal energy is being taken in from the surroundings (1)
Worked solution
Sodium hydroxide's temperature rises by 32.5 − 20.0 = 12.5°C, while ammonium chloride's temperature falls by 20.0 − 14.5 = 5.5°C. A temperature rise means thermal energy has been released into the water, so dissolving sodium hydroxide is exothermic. A temperature fall means thermal energy has been absorbed from the water, so dissolving ammonium chloride is endothermic. This shows that dissolving is not always the same type of process — it depends on the particular substance being dissolved.
Question 29
SupplementExothermic and endothermic reactions[5]
Ethanol burns in a spirit burner:
C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(g)
This reaction is exothermic overall, even though energy must first be put in to break the bonds in the reactants.
(a)
State whether breaking bonds is an exothermic or endothermic process, and whether making bonds is an exothermic or endothermic process.
[2]
(b)
Explain, in terms of bond breaking and bond making, why the combustion of ethanol is exothermic overall even though bonds must first be broken.
energy must be absorbed to break the bonds in the reactants (ethanol and oxygen) (1)
energy is released when the new bonds in the products (carbon dioxide and water) form (1)
because more energy is released forming the new bonds than is absorbed breaking the old ones, the overall reaction releases energy / is exothermic (1)
Worked solution
Breaking any chemical bond always requires an input of energy (endothermic), while forming any new bond always releases energy (exothermic). When ethanol burns, energy must first be absorbed to break the C–C, C–H, C–O, O–H and O=O bonds in the reactants. Then energy is released as new C=O and O–H bonds form in the carbon dioxide and water produced. Combustion reactions like this are exothermic overall because the amount of energy released forming the new, more stable bonds in the products is greater than the amount of energy absorbed breaking the bonds in the reactants — the surplus energy is released to the surroundings as heat.
Question 30
SupplementExothermic and endothermic reactions[8]
Methane burns completely in a plentiful supply of oxygen:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
Use the bond energies in the table to answer the questions below.
Bond
Bond energy (kJ/mol)
C–H
413
O=O
496
C=O
805
O–H
463
(a)
Calculate the total energy needed to break all the bonds in the reactants.
[2]
(b)
Calculate the total energy released when all the bonds in the products form.
[2]
(c)
Calculate ∆H for the complete combustion of methane, showing your working, and state whether the reaction is exothermic or endothermic.
[3]
(d)
State whether the products would be drawn at a higher or lower energy level than the reactants on a reaction pathway diagram for this reaction, and explain your answer.
∆H = energy to break bonds − energy released forming bonds (1)
∆H = 2644 − 3462 = −818 kJ/mol (1)
exothermic (1)
ECF from (a) and (b)
(d)[1]
lower; because the reaction is exothermic, the products have less energy than the reactants
ECF from (c)
Worked solution
Breaking the four C–H bonds in methane and the two O=O bonds in oxygen needs (4 × 413) + (2 × 496) = 2644 kJ. The products are one CO2 molecule (with 2 C=O bonds) and two H2O molecules (with 4 O–H bonds in total), so forming these releases (2 × 805) + (4 × 463) = 3462 kJ. This gives ∆H = 2644 − 3462 = −818 kJ/mol, confirming that methane combustion is strongly exothermic, as expected for a fuel used for heating. On a reaction pathway diagram, this means the products would be drawn at a lower energy level than the reactants, with the large drop in energy between them representing the ∆H value released to the surroundings.