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Enthalpy is what we call the amount of heat in a chemical system. When a chemical reaction takes place, energy in the form of heat gets absorbed or evolved. This transfer of heat energy changes the total enthalpy of a system. When a reaction absorbs energy, the total enthalpy increases, but when it releases energy, the total enthalpy decreases. We cannot directly measure the total enthalpy in a system, but we can measure the enthalpy change (∆H) that takes place in a chemical reaction.
Enthalpy change gives us a good idea about whether a reaction will take place or not. Even better, it helps us put energy to good use.
When heat energy flows in or out of a chemical reaction at constant pressure, enthalpy change takes place.
Enthalpy change (∆H) is the amount of heat energy transferred in a chemical system at constant pressure.
We measure enthalpy change using a method known as calorimetry. This involves observing the temperature change when a system absorbs or releases heat energy.
Essentially, we record the temperature change in a known amount of substance (generally water) with a known specific heat. Enthalpy change happens when the chemical reaction under study absorbs or releases heat (as shown in the diagram below).
A simple calorimeter, open textbc.ca
You do not need to carry out calorimetry every time you want to know the enthalpy change of a reaction. Later in this article, you will discover standard enthalpy change, but before that, let us look at endothermic and exothermic reactions.
Reactions can either be exothermic or endothermic.
Endothermic reactions absorb heat energy from the surroundings.
Exothermic reactions release heat energy into the surroundings.
Photosynthesis is an example of an endothermic reaction in which plants take in energy from the sun. On the other hand, respiration is an exothermic reaction that releases the energy stored in glucose.
We can illustrate exothermic and endothermic reactions using reaction profiles or enthalpy diagrams. They show the enthalpy levels of the reactants and products in an exothermic and endothermic reaction.
Enthalpy diagram for an exothermic and endothermic reaction, Olive [Odagbu] StudySmarter Originals
The y-axis shows enthalpy, while the extent of the reaction is shown on the x-axis. The dotted lines show the enthalpy levels of the reactants and products. The difference between the levels is the enthalpy change (∆H) as the reactants turn into the products. What differences can you see between endothermic and exothermic reactions?
In endothermic reactions, the products have a higher enthalpy than the reactants.
In exothermic reactions, the reactants have a higher enthalpy than the products.
Endothermic reactions have a positive ∆H, while exothermic reactions have a negative ∆H.
We can also use reaction profiles to graph a reaction pathway, which is a line that shows how the reactants turn into the products in terms of enthalpy. You draw a reaction pathway on similar axes to the ones we used in the enthalpy diagrams above. However, instead of straight upwards or downward arrows, we use a curved line to show the enthalpy change as the bonds break in the reactants and new bonds form in the products.
Here is an example of a reaction pathway. You can see that the products have a lower enthalpy than the reactants, meaning this is an exothermic reaction. What else do you notice?
Exothermic reaction profile, Olive [Odagbu] StudySmarter Originals
Activation energy is the minimum amount of energy required for a reaction to take place.
The reaction pathway of an endothermic reaction looks like a mirror image of the reaction pathway of an exothermic reaction. Here, the reactants have lower enthalpy than the products.
Endothermic reaction profile, Olive [Odagbu] StudySmarter Originals
Earlier, we said that you do not need to carry out calorimetry to know the enthalpy change of a reaction. Fortunately, scientists have done much of the hard work for us. That means we can use values called standard enthalpy changes (∆Hϴ) to predict the enthalpy change of a reaction.
Standard enthalpy changes are reactions done under standard conditions with the reactants and products in their standard states. What are standard conditions? A simple reaction can give different results if the reactants and products are in different states or under high pressure. For chemists to know they experimented in the same environment, they use standard conditions. These are a set of criteria, which are:
298 K or 25 ºC.
Pressure at 1 bar or 100kPa.
In the case of solutions, 1 mol dm-3 concentration.
What about standard states? Under standard conditions, some substances are liquid, while others are gas, and others are solid. The state of a chemical under standard conditions is its standard state. For example, under the standard conditions of 298 K and 100 kPa, oxygen is a gas (g) - and so its standard state is gaseous.
We usually give enthalpy change values as standard values, which we call standard enthalpy changes.
You need to know about a few different types of standard enthalpy changes. In Hess’ Law, you can learn how these enthalpy changes help us to calculate ∆H for reactions we can’t measure with calorimetry. The enthalpy changes we will consider here are:
Standard enthalpy change of reaction.
Standard enthalpy change of combustion.
Standard enthalpy change of formation.
The enthalpy change of reaction (∆Hϴr) is the enthalpy change when reactants react to make products as stated in the chemical equation under standard conditions and in their standard states.
Consider the following equation where calcium oxide reacts with water. What does it mean?
CaO(s) + H2O(l) → Ca(OH)2(s) ∆Hϴr = -63.7 kJ·mol-1
The standard enthalpy change of reaction when 1 mole of calcium oxide reacts with 1 mole of water to make 1 mole of calcium hydroxide is -63.7 kilojoules per mole. 63.7 kJ·mol-1 is the enthalpy change that occurs when this reaction takes place. That means the system’s enthalpy decreases by 63.7 kJmol-1.
Notice also that we show the standard states of the reactants and products with state symbols after each substance, i.e., (s) for solid and (l) for liquid.
Usually, when we write the enthalpy change of reaction (∆Hr), we leave out the r. We just assume that ∆H means the enthalpy change of reaction.
The standard enthalpy change of formation (∆fHϴ) is the enthalpy change when one mole of a compound forms from its elements under standard conditions in their standard states.
The equation below shows the standard enthalpy change of formation for calcium carbonate.
Ca(s) + C(s) + O2(g) → CaCO3(s) ∆fHϴ = 52.5 kJmol-1
You will notice that the equation produces only 1 mole of calcium carbonate. You will also notice that we used (or
) moles of oxygen in the equation. While this might seem odd, you must get only 1 mole of a compound when you write an equation for the standard enthalpy of formation. In this case, it’s perfectly acceptable to use fractions to balance equations for ∆fHϴ.
The ∆fHϴ value for any element in its standard state is 0. Since we cannot know the absolute enthalpy, we use 0 as the baseline of an element’s most stable form at the standard state. We then measure the enthalpy of formation values for compounds from that 0 baseline.
We know the standard enthalpies of formation (∆fHϴ) for many compounds. You can find them on an enthalpy table in your exam. You put these values into the equation for Hess’s Law when calculating the enthalpy change for a reaction.
The standard enthalpy change of combustion (∆cHϴ) is the enthalpy change when 1 mole of a substance burns completely in oxygen under standard conditions with all the reactants and products in their standard states.
Enthalpies of combustion will always be negative since combustion is an exothermic reaction. Take, for instance, the equation for the standard enthalpy of combustion of methane shown below:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ∆cHϴ = -890 kJmol-1
Sometimes, you may notice similarities between equations for ∆fHϴ and ∆cHϴ. Have a look at the following two equations. What differences can you spot?
C(s) + O2(g) → CO2(g) ∆fHϴ = -393.5 kJmol-1
C(s) + O2(g) → CO2(g) ∆cHϴ = -393.5 kJmol-1
There are no differences. Both of these equations are the same, except one is for the standard enthalpy of combustion of carbon, while the other is the standard enthalpy of formation of carbon dioxide. This can happen in some simple cases. Another example would be the standard enthalpy of the formation of water and the standard enthalpy of combustion of hydrogen, as shown below.
H2(g) + ½O2(g) → H2O(l) ∆fHϴ = -285.8 kJmol-1
H2(g) + ½O2(g) → H2O(l) ∆cHϴ = -285.8 kJmol-1
We will now consider the enthalpy change of neutralisation.
When an acid reacts with an alkali, we call the enthalpy change that takes place the enthalpy change of neutralisation.
The standard enthalpy change of neutralisation (∆nHϴ) is the enthalpy change when an acid solution and an alkali solution react under standard conditions to form 1 mole of water.
Under standard conditions, all reactions between a strong acid and a strong base have similar values for the standard enthalpy of neutralisation. This is always a negative value, usually between -57 and -58 kJmol-1. Essentially, reactions that involve a strong acid and a strong base are just reactions between hydrogen ions and hydroxide ions. The other ions are spectator ions. The values are alike because acids and bases always react similarly: hydrogen and hydroxide ions react to form water.
If you like to read more about this, check out the explanation on Acids and Bases.
Weak acids and bases do not fully dissociate in water. In reactions that involve weak acids and bases, some of the energy changes involved get used to dissociate the weak acid or base. This means that not as much energy gets evolved to make hydrogen and hydroxide ions. So, in reactions with weak acids or bases, the standard enthalpies of neutralisation are always a little less exothermic.
Previously, we said that we use calorimetry to figure out the enthalpy change of a reaction. Let us examine how this works.
Essentially, we measure the change in temperature in the surroundings when a reaction releases or absorbs heat. We use a thermometer to measure the temperature change. Then we convert the temperature change (∆T) to heat energy (q) with the equation:
q = m c ∆T
Here:
We can measure the temperature change in either Kelvin or ºC. That is because a change of 1 Kelvin is the same as a change of 1ºC.
You might have noticed that the equation is for q, the heat energy released or evolved, not ∆H. That is okay since we can convert joules to kilojoules per mole by first dividing by the number of moles and then by 1000.
Here is an example of how you might use q = mc∆T to figure out the enthalpy change per mole of substances that are soluble in water.
Shane dissolved 0.068 moles of potassium iodide in 48g of water. He observed a decrease in temperature of 14.8ºC.
Calculate the enthalpy change in kJmol-1. Water has a specific heat capacity of 4.48 Jg-1 k-1.
q = mc∆T
q = (48g)(4.48 Jg-1 k-1)(14.8ºC)
q = 3182.592 Joules
To find the energy per mole of potassium iodide, divide by 0.068 moles of KI used.
Divide by 1000 to convert to kJmol-1.
Place a ‘+’ sign before the enthalpy change value because the reaction took heat from the surroundings. The temperature of the water decreased.
Check out other examples of this formula for enthalpy change in the explanation of Calorimetry.
Since enthalpy change is a state function, we can use the following equation to calculate the enthalpy change of formation of a substance:
∆Hϴ = ∆fHϴ(products) - ∆fHϴ(reactants)
You can use calorimetry and the equation q = mc∆T to calculate the enthalpy change. You could also use an enthalpy cycle or the equation ΔH = enthalpy change for broken bonds + enthalpy change for formed bonds.
You can calculate the standard enthalpy by conducting calorimetry under standard conditions with the reactants and products in their standard states. Scientists have done much of the hard work for us, so we don’t have to conduct a calorimetry experiment every time we want to know the enthalpy change of a reaction. You can simply look up standard enthalpy change values for many substances in an enthalpy table.
You can calculate the enthalpy change of combustion by using the enthalpies of formation of the products and reactants in the equation:
∆cHϴ = ∑∆fHϴ(Reactants) - ∑∆fHϴ(Products).
Enthalpy change (∆H) is the amount of heat energy transferred in a chemical system at constant pressure.
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