Energy in reactions

You’re being asked which condition must be kept fixed when we define the heat energy change of a system as enthalpy change ($\Delta H$). This is a core thermochemistry definition used throughout IB Chemistry energetics topics.

In IB Chemistry, enthalpy ($H$) is described as the heat content of a system. Since we can’t measure an absolute value of $H$ directly, we focus on changes, $\Delta H$.

By definition, $\Delta H$ is the heat transferred between a system and surroundings when the pressure is kept constant. This relationship is written as:

$q_p=\Delta H$

where $q_p$ means the heat exchange measured at constant pressure.

Each incorrect option fails because it describes something that may affect measurements, but isn’t the defining condition for $\Delta H$:

• Concentration: Changing concentration can change how much heat is produced overall, but it is not the condition used to define enthalpy change.
• Volume: If heat is measured at constant volume, it corresponds to a change in internal energy ($\Delta U$), not enthalpy.
• Temperature: Temperature often changes during reactions, and “constant temperature” would describe an isothermal process; it is not part of the definition of $\Delta H$.

Because enthalpy change is defined as heat energy change at constant pressure, the correct answer is pressure.

In thermodynamics, enthalpy change ($\Delta H$) tracks energy changes for a system in terms of heat ($q$) exchanged with the surroundings. To answer the question, we identify which condition must be held constant for heat to be described as “enthalpy change.”

Concentration describes how much solute is present in a given volume of solution. It can influence reaction rate or the total heat produced, but it is not the fixed condition used in the definition of enthalpy. So this option can be eliminated.

At constant volume (for example, in a rigid bomb calorimeter), the system cannot expand, so expansion work ($P\Delta V$) is not done. Under those conditions, the measured heat corresponds to the change in internal energy ($\Delta U$), not $\Delta H$. Therefore, volume is not correct.

We often use temperature change to calculate heat (for instance with $q=mc\Delta T$), but keeping temperature constant would describe an isothermal situation. Enthalpy depends on temperature, but the defining condition for $\Delta H$ is not “constant temperature.”

Enthalpy is defined by:

$H=U+PV$

When pressure is held constant, the enthalpy change equals the heat transferred at constant pressure:

$\Delta H=q_p$

This is why many school experiments done in open containers (at atmospheric pressure) are treated as measurements of enthalpy change.

To decide whether the reaction releases or absorbs energy, start by pulling out the enthalpy ($H$) values stated for each side of the reaction.

• Enthalpy of reactants ($H_{\text{reactants}}$) = $400{\text{ kJ mol}}^{-1}$
• Enthalpy of products ($H_{\text{products}}$) = $150{\text{ kJ mol}}^{-1}$

Enthalpy change, $\Delta H$, is calculated by subtracting the reactants’ enthalpy from the products’ enthalpy. This is the key relationship used for Energy changes in a reaction.

$\Delta H=H_{\text{products}}-H_{\text{reactants}}$

Put the given enthalpy values into the formula:

$\begin{array}{rl}\Delta H& =150{\text{ kJ mol}}^{-1}-400{\text{ kJ mol}}^{-1}\\ \Delta H& =-250{\text{ kJ mol}}^{-1}\end{array}$

The sign of $\Delta H$ tells you what happened to energy:

• If $\Delta H$ is negative ($\Delta H<0$), energy left the system and went to the surroundings, so the reaction is exothermic.
• If $\Delta H$ is positive ($\Delta H>0$), energy entered the system from the surroundings, so the reaction is endothermic.

Because $\Delta H=-250{\text{ kJ mol}}^{-1}$, this reaction must be exothermic.

Compare each option to what we found:

• Options A and B have the wrong sign/value (they use $+250$ instead of $-250$).
• Option C has the correct number but the wrong reaction type (it says endothermic).

The calculation gives $\Delta H=-250{\text{ kJ mol}}^{-1}$, and the negative value means energy is released, so the reaction is exothermic.

Start by noting the energy levels that the question gives for each side.

• Enthalpy of reactants ($H_{\text{reactants}}$) = $400{\text{ kJ mol}}^{-1}$
• Enthalpy of products ($H_{\text{products}}$) = $150{\text{ kJ mol}}^{-1}$

With the Enthalpy Level Diagram Approach, imagine enthalpy on the vertical axis:

1. Mark a horizontal level at $400$ for the reactants.
2. Mark a horizontal level at $150$ for the products.

Because the products sit lower than the reactants, the diagram shows a drop in energy. That drop means energy has been transferred out to the surroundings.

The enthalpy change is still found using:

$\Delta H=H_{\text{products}}-H_{\text{reactants}}$

Substitute the values:

$\begin{array}{rl}\Delta H& =150-400\\ \Delta H& =-250{\text{ kJ mol}}^{-1}\end{array}$

The minus sign matches the downward step on the diagram.

From both the sketch and the calculation:

• Products lower than reactants means energy is released, so the reaction is exothermic.
• The enthalpy change is $\Delta H=-250{\text{ kJ mol}}^{-1}$.

Elimination based on what the diagram shows:

• Options A and C are wrong because they call it “Endothermic,” but the energy level went down.
• Option B is wrong because it uses $+250$, which would suggest energy was taken in.

Since the products are at a lower enthalpy than the reactants, the reaction releases energy and has a negative enthalpy change: exothermic with $\Delta H=-250{\text{ kJ mol}}^{-1}$.

To classify the reaction, compare the enthalpy of the starting substances to the enthalpy of the substances formed.

We are told the enthalpy of the reactants is $400{\text{ kJ mol}}^{-1}$ and the enthalpy of the products is $150{\text{ kJ mol}}^{-1}$. Because the products end up with less energy, the “missing” energy must have been released to the surroundings as the reaction happened.

If energy is released overall, the reaction is exothermic. For an endothermic change, the products would need to be higher in enthalpy than the reactants.

So any choice that describes the reaction as 'Endothermic' cannot be correct here.

Now calculate the enthalpy change using the rule: products minus reactants.

$\Delta H=H_{\text{products}}-H_{\text{reactants}}$

Substituting the values: $\Delta H=150-400=-250{\text{ kJ mol}}^{-1}$

A negative $\Delta H$ always goes with an exothermic reaction.

Option B claims “exothermic” but shows a positive $\Delta H$, which doesn’t make sense. That leaves Option D as the only choice that matches both the sign and the reaction type.

We answer by remembering the definition of activation energy ($E_{\text{a}}$). In collision theory, it is the smallest amount of energy particles must have when they collide for a reaction to happen.

On an energy profile diagram, this corresponds to the energy barrier that must be climbed to reach the transition state (the highest point of the curve) starting from the energy level of the reactants.

This choice is describing the overall energy change from the beginning to the end of the reaction. That measurement is the enthalpy change ($\Delta H$), not the activation energy, because it compares reactants directly to products rather than to the peak.

A measurement that starts at the products and goes up to the peak is describing how much energy would be needed for the reaction to run backwards.

For an exothermic reaction, products sit at a lower energy than reactants, so this products-to-peak distance would be larger than the forward activation energy.

Energy profile diagrams are usually about relative energy levels. Using a “zero energy” baseline is not how activation energy is defined on these graphs.

$E_{\text{a}}$ is the extra energy needed to raise the reactants up to the transition state, not the total energy above some absolute reference.

Activation energy is the “uphill climb” from where the reactants begin to the highest point of the curve. So, on the energy profile, it is the vertical distance from the reactants’ energy level to the peak.

Imagine the standard energy profile for an exothermic reaction.

The vertical axis shows Potential Energy, and the horizontal axis shows Reaction Progress. Because the reaction is exothermic, the reactants start at a higher energy level and the products finish at a lower energy level.

Reactions need energy at the start to disrupt existing bonds so new bonds can form. On the diagram this appears as a rise to a maximum.

That maximum point is the Transition State (also called the activated complex). It is the point of greatest energy the system must reach for the reaction to proceed.

The activation energy ($E_{\text{a}}$) is the minimum energy needed for the reaction to happen.

On the diagram, that is the energy increase required to go from the reactants’ starting level up to the top of the barrier (the peak). So the $E_{\text{a}}$ arrow begins at the reactant energy level and points vertically to the peak.

Using this picture, we can reject the incorrect measurements:

• Option A: products to peak corresponds to activation energy for the reaction running in reverse.
• Option C: reactants to products is the enthalpy change, $\Delta H$.
• Option D: zero to peak is not appropriate because the reactants already have energy; we measure the extra energy needed from their level.

From the visual model, $E_{\text{a}}$ is the vertical distance from the reactants’ energy level up to the peak of the curve.

For an exothermic energy profile, the reactants begin at a higher energy than the products. The reaction still needs to get over an energy barrier called the activation energy ($E_{\text{a}}$).

We can identify the correct distance by checking what each incorrect option is actually measuring.

The vertical difference between reactants and products is the enthalpy change ($\Delta H$).

Although $\Delta H$ is negative for an exothermic reaction (energy released), it does not describe the energy needed to get the reaction started. So this option is not $E_{\text{a}}$.

A distance measured from the products up to the peak is the energy barrier for the reverse reaction.

Because we want the activation energy for the forward reaction, this is the wrong measurement.

Starting from a “zero energy level” and measuring to the peak is not meaningful here. Activation energy is a relative quantity: it is the additional energy reactant particles must gain to reach the transition state.

Using zero ignores the fact that reactant molecules already contain internal energy.

After eliminating the other meanings, the remaining distance is the standard one: the “hill” from the reactants’ energy level up to the peak (transition state). That is the activation energy.

The prompt tells you photosynthesis is an endothermic process. In chemistry, that means the reacting system (here, the plant’s cells) has to take in energy from the surroundings (sunlight) so the reaction can happen.

This links directly to Topic 3: Types of chemical reaction, especially how endothermic and exothermic reactions are described by energy transfer.

For an endothermic reaction, the products store more energy than the reactants, so the enthalpy (energy content) is higher for products than reactants ($H_{products}>H_{reactants}$).

In photosynthesis, $C{O}_2$ and $H_{2}O$ start at a lower energy level. Building glucose ($C_6{H}_{12}{O}_6$) creates a higher-energy product, so energy has to be supplied continuously. Light energy is what provides that input, effectively raising the reactants up to the higher-energy products.

Check the first two choices:

• Option A focuses on cooling to help water uptake. Plants can lose heat through processes like transpiration, but that is not the main chemical role of light in photosynthesis.
• Option B says light is a catalyst. Catalysts lower activation energy ($E_a$) and speed up reactions, but they don’t serve as the energy source that makes products higher in energy. In photosynthesis, light energy is taken in as part of the energy requirement, not acting like a catalyst.

Option C suggests light energy is produced when glucose forms. That would match an exothermic situation where energy is released. But photosynthesis is stated to be endothermic, meaning energy must be absorbed. If it released light while making glucose, it wouldn’t need incoming sunlight to proceed.

Because endothermic reactions require an energy input so the products can exist at a higher stored-energy level than the reactants, only Option D matches the thermodynamic description.

In IB MYP Chemistry, we treat reactions as processes that involve energy changes. An endothermic reaction is one where energy is absorbed from the surroundings. For photosynthesis, plants take in light energy to change carbon dioxide and water into glucose and oxygen.

$6C{O}_2+6H_{2}O+\text{light energy}\to C_6{H}_{12}{O}_6+6O_2$

The question is asking what role that light energy plays in this energy-absorbing reaction.

Option A claims the purpose of light is cooling. That’s not correct. When a leaf absorbs light, it often warms up. Plants may cool themselves by transpiration (water evaporating from leaves), but that is a different process and not the main role of light in the photosynthesis reaction itself.

Option B treats light as a catalyst. A catalyst (often an enzyme in biology) speeds up a reaction by lowering activation energy and is not used up. Light is not a substance and isn’t a catalyst; it is energy. In photosynthesis, enzymes help the steps occur, but the light provides the energy input that drives the reaction forward.

Option C says light is released, which would describe an exothermic reaction. But the question clearly states photosynthesis is endothermic, so energy must go in rather than come out. The chemical products are glucose and oxygen; light energy is the required input.

After removing the incorrect ideas, the correct description is the one that fits endothermic reactions: the products (glucose) store more chemical potential energy than the reactants ($C{O}_2$ and $H_{2}O$). Light energy must be absorbed to supply the energy needed to reach that higher-energy product state.

Photosynthesis is the life process in which autotrophs (such as plants) use chlorophyll to take in energy from sunlight. That captured energy allows them to build organic molecules—especially glucose—from carbon dioxide and water.

Looking at the energy change, photosynthesis is a typical endothermic reaction because energy is taken in from the surroundings and stored as new chemical bonds form.

In photosynthesis, energy moves in a clear pathway:

1. Input: Sunlight (light energy) is absorbed by pigments.
2. Change: That light energy powers a set of chemical reactions.
3. Result: Energy ends up stored in the bonds of glucose ($C_6{H}_{12}{O}_6$).

So overall, photosynthesis changes light energy into chemical energy.

Option C claims chemical energy becomes light energy. That describes situations where light is produced from chemicals—like bioluminescence or burning fuels to make light—not what plants do to make food. Photosynthesis mainly takes in light rather than producing it as the key energy output.

Option B says the products have lower potential energy than the reactants. When products are lower in energy, the reaction is exothermic because energy is released. But photosynthesis absorbs energy and locks it into glucose, meaning the products must have more stored (potential) energy than the reactants.

Even though real processes can waste some energy as heat, photosynthesis is not best described as mainly giving off thermal energy. The central idea is that sunlight is taken in and conserved as chemical energy in newly formed molecules.

Using the core definition of photosynthesis, the correct description is that sunlight is absorbed and then stored as chemical energy in organic molecules.

To answer this, focus on what photosynthesis does: plants use sunlight to turn carbon dioxide and water into glucose ($C_6{H}_{12}{O}_6$) and oxygen. This is an endothermic process because it needs a continuous energy input from the surroundings. This links to Topic: Endothermic and exothermic reactions.

Option C says chemical energy changes into light energy. That fits examples like fireflies (bioluminescence) or a burning candle (combustion). In photosynthesis, the plant’s role is to take in light energy to drive reactions, not to generate light as the main output.

Option D focuses on heat being released. Giving energy to the surroundings is typical of exothermic reactions (for example, burning wood). Photosynthesis is endothermic, so it primarily requires the absorption of solar energy rather than releasing thermal energy to the air.

Because photosynthesis absorbs energy, that energy ends up stored in the chemical bonds of glucose. This means the products must have higher potential energy than the reactants ($C{O}_2$ and $H_{2}O$). Option B states the opposite, so it cannot be correct.

After eliminating the others, Option A is left. Photosynthesis takes solar energy and converts it into chemical energy stored in the bonds of glucose molecules, which is the best description of the energy flow.

An endothermic reaction is defined as a reaction where energy is moved from the surroundings into the reacting chemicals (the system).

So in an endothermic change, the system takes in heat from the surroundings, which often makes the surroundings feel cooler or causes their temperature to drop.

Enthalpy ($H$) is the heat energy stored in a system. The enthalpy change is found using:

$\Delta H=H_{\text{products}}-H_{\text{reactants}}$

In an endothermic reaction, energy is absorbed, meaning the products end up with more stored energy than the reactants ($H_{\text{products}}>H_{\text{reactants}}$). That makes $\Delta H$ positive.

Using the definition and the sign of $\Delta H$, we can remove the wrong statements:

• Option B and Option D are wrong because they describe heat leaving the system, which is what happens in an exothermic reaction.
• Option C is wrong because $\Delta H$ being negative means energy is released overall (exothermic), not absorbed.

Only one option correctly states both key features: heat is absorbed from the surroundings and the enthalpy change is positive.

Start with the word endothermic: “endo-” suggests “into,” and “-thermic” relates to heat. That points to heat moving into the system during the reaction.

This idea is central in Topic: Types of chemical reaction, especially the section on Endothermic and exothermic reactions.

Think about an enthalpy (energy) profile diagram. For an endothermic reaction, the products sit at a higher enthalpy level than the reactants.

Imagine:

1. A lower horizontal line for Reactants.
2. A higher horizontal line for Products.

Moving from the lower line to the higher one requires an input of energy, so the system must absorb energy from outside.

The enthalpy change, $\Delta H$, is calculated by subtracting the reactants’ enthalpy from the products’ enthalpy:

$\Delta H=H_{products}-H_{reactants}$

Because the products are higher than the reactants on the diagram ($H_{products}>H_{reactants}$), the difference must be positive.

$\Delta H>0$

Now use the diagram idea to reject statements that match exothermic reactions. Exothermic reactions have $\Delta H$ negative and give out heat.

Options C and D do not fit because they say $\Delta H$ is negative.

If $\Delta H$ is positive, the system ends with extra energy in the products. That extra energy must come from the surroundings.

So the system absorbs heat from the surroundings, which rules out Option B because it claims heat is released.

The energy profile confirms the story: products higher than reactants means $\Delta H$ is positive, and the system must take in heat from the surroundings overall.

Begin with the prefix 'endo-', which means “inside” or “into.”

In chemistry, an endothermic reaction is one where thermal energy is taken in by the system from the surroundings. This is part of Topic 3: Types of chemical reaction, under Endothermic and exothermic reactions.

Because endothermic reactions require heat to enter the system, any option that says the system is 'releasing' heat must be incorrect. Releasing heat describes an exothermic reaction.

So we can eliminate Option B and Option D.

Next, decide the sign of $\Delta H$.

Using the IB sign convention: energy going into the system is positive ($+$), and energy going out is negative ($-$). Since an endothermic reaction absorbs energy, the system’s enthalpy increases.

Therefore, $\Delta H$ must be positive, so Option C (negative $\Delta H$) is incorrect.

After removing the choices that contradict endothermic heat flow and the sign convention, the remaining statement is the one that correctly describes an endothermic reaction.