Endothermic and Exothermic Reactions - AP Chemistry

ΔH is always positive for an endothermic reaction, and ΔG is always negative for a spontaneous reaction. Given the equation delta G = ΔH – T(ΔS), T(ΔS) is positive, so ΔS is positive. We do not know anything about the equilibrium of the reaction.

In an endothermic reaction, heat is absorbed and converted into chemical energy. Heat is the sum of molecular kinetic energy in a sample, so kinetic energy is converted into chemical energy. Note that temperature is a measure of heat, and is not a type of energy in itself.

The rate of a reaction will always increase with temperature.

Do not confuse the rate at which a reaction proceeds with the final equilibrium of the reaction. Yes, increasing the temperature of an exothermic reaction will push the equilibrium towards the reactant's side, but this equilibrium will be achieved much more quickly at higher temperatures. Kinetics and thermodynamics are two completely different ways to observe a reaction.

In exothermic and endothermic reactions, heat can be considered to be either a product or reactant, respectively. Increasing the temperature of the an exothermic reaction is effectively the same as adding heat as a product, causing the reaction to shift to the left. This will cause the reverse reaction to increase in rate.

From the question stem, we are told that the breakdown of 2 moles of gaseous requires an absorption of of energy. For simplification, we can write out the reaction as:

By definition, the heat of formation of a compound is the enthalpy change that occurs for the formation () of that compound from its elements in their standard state. Since and are in their standard state, we know that they have a heat of formation of zero, thus simplifying the calculation. Thus, if we reverse the above reaction, we see that:

Notice that by reversing the reaction, we are also reversing the sign for the enthalpy of the reaction. Now, to calculate the of as kilojoules per mole, we need to multiply the reaction by one-half to obtain:

We can calculate the standard change in enthalpy for the reaction using the following equation:

of

of

of

Plug in known values and solve.

A negative indicates that the reaction is exothermic. This means that the reaction will release energy (usually in the form of heat) to the surroundings.

A positive indicates that the reaction is endothermic. This means that the reaction will absorb energy (usually in the form of heat) from the surroundings.

Therefore, this overall reaction is endothermic.

Negative enthalpy change() indicates that heat is on the product side of the reaction, or, is released by the reaction. This is also the definition of an exothermic reaction.

Endergonic and exergonic are terms used to describe the free energy of a reaction, so they do not apply in this case since there is not enough information to determine the free energy produced by this reaction without knowing the temperature (delta G = delta H - T delta S). The negative enthalpy allows us to definitively say that the reaction is exothermic.

ΔH is always positive for an endothermic reaction, and ΔG is always negative for a spontaneous reaction. Given the equation delta G = ΔH – T(ΔS), T(ΔS) is positive, so ΔS is positive. We do not know anything about the equilibrium of the reaction.

In an endothermic reaction, heat is absorbed and converted into chemical energy. Heat is the sum of molecular kinetic energy in a sample, so kinetic energy is converted into chemical energy. Note that temperature is a measure of heat, and is not a type of energy in itself.

The rate of a reaction will always increase with temperature.

Do not confuse the rate at which a reaction proceeds with the final equilibrium of the reaction. Yes, increasing the temperature of an exothermic reaction will push the equilibrium towards the reactant's side, but this equilibrium will be achieved much more quickly at higher temperatures. Kinetics and thermodynamics are two completely different ways to observe a reaction.

In exothermic and endothermic reactions, heat can be considered to be either a product or reactant, respectively. Increasing the temperature of the an exothermic reaction is effectively the same as adding heat as a product, causing the reaction to shift to the left. This will cause the reverse reaction to increase in rate.

From the question stem, we are told that the breakdown of 2 moles of gaseous requires an absorption of of energy. For simplification, we can write out the reaction as:

By definition, the heat of formation of a compound is the enthalpy change that occurs for the formation () of that compound from its elements in their standard state. Since and are in their standard state, we know that they have a heat of formation of zero, thus simplifying the calculation. Thus, if we reverse the above reaction, we see that:

Notice that by reversing the reaction, we are also reversing the sign for the enthalpy of the reaction. Now, to calculate the of as kilojoules per mole, we need to multiply the reaction by one-half to obtain:

We can calculate the standard change in enthalpy for the reaction using the following equation:

of

of

of

Plug in known values and solve.

A negative indicates that the reaction is exothermic. This means that the reaction will release energy (usually in the form of heat) to the surroundings.

A positive indicates that the reaction is endothermic. This means that the reaction will absorb energy (usually in the form of heat) from the surroundings.

Therefore, this overall reaction is endothermic.

Negative enthalpy change() indicates that heat is on the product side of the reaction, or, is released by the reaction. This is also the definition of an exothermic reaction.

Endergonic and exergonic are terms used to describe the free energy of a reaction, so they do not apply in this case since there is not enough information to determine the free energy produced by this reaction without knowing the temperature (delta G = delta H - T delta S). The negative enthalpy allows us to definitively say that the reaction is exothermic.

ΔH is always positive for an endothermic reaction, and ΔG is always negative for a spontaneous reaction. Given the equation delta G = ΔH – T(ΔS), T(ΔS) is positive, so ΔS is positive. We do not know anything about the equilibrium of the reaction.

In an endothermic reaction, heat is absorbed and converted into chemical energy. Heat is the sum of molecular kinetic energy in a sample, so kinetic energy is converted into chemical energy. Note that temperature is a measure of heat, and is not a type of energy in itself.

The rate of a reaction will always increase with temperature.

Do not confuse the rate at which a reaction proceeds with the final equilibrium of the reaction. Yes, increasing the temperature of an exothermic reaction will push the equilibrium towards the reactant's side, but this equilibrium will be achieved much more quickly at higher temperatures. Kinetics and thermodynamics are two completely different ways to observe a reaction.