Classwork Series and Exercises {Chemistry – SS2}: Energy Changes In Chemical Reactions

Chemistry SS 2 Week 2

Topic: Energy Changes in Chemical Reactions

Introduction

Energy changes take place during the chemical reactions. Almost all the chemical reactions involve some sort of change in energy; either they absorb energy or release energy, generally in the form of heat.

Heat is a form of energy that always flows from an object at high temperature to an object at lower temperature, but not other way around. For instance, if you open the front door of your house in the winter time, the heat flows out but not vice verse. When you place a piece of ice on the table top, it begins to melt because the heat flows from surrounding into the ice thereby producing the melting process.

In chemistry, the energy changes in chemical reactions are studied by a special branch of chemistry known as thermochemistry, which is defined as the study of heat changes in chemical reactions.

To understand energy changes, first we must define two things, the system and the surrounding. The system is a specific part of the universe on which we concentrate. In chemistry for example, if you are carrying out an experiment in a beaker, the reaction mixture and the beaker constitute the system. The rest of the universe outside the system is considered as surrounding.

There are three types of systems:

An open system: An open system is the system that can exchange mass and energy (heat) with the surrounding, for example heating water in an open beaker. Let us set up a simple experiment, take a certain quantity of water in an open beaker (open system) and start to heat it up. As the water begins to boil, some water converts into vapor and escapes to the surrounding above the beaker (loss of mass). At the same time, the heat also escapes to the surrounding through the opening as well as through the walls of the beaker.

A closed system: A closed system is the system where only the heat is exchanged with the surrounding. We take the same beaker and close the lid and try to heat it up (closed system). Since the lid is closed, the mass cannot escape. However, the heat can still escape to the surrounding through the lid and through the walls of the beaker.

An isolated system: An isolated system is the system where neither heat nor mass is exchanged with the surrounding. We completely insulate the beaker and try to heat it up (closed system). Since it is completely insulated, neither mass nor heat can escape to the surrounding.

In terms of the energy changes that take place during chemical reactions, a reaction may be either exothermic or endothermic, terms which were first coined by the French Chemist Marcellin Berthelot (1827 − 1907).

Exothermic Reactions

Exothermic chemical reactions release energy. The released energy may be in the form of heat, light, electricity, sound or shock waves either singly or in combinations.

A few examples of exothermic reactions are:

• Mixing of acids and alkalis (releases heat)
• Combustion of fuels (releases heat and light)

Endothermic Reactions

Endothermic chemical reactions absorb energy. The energy absorbed may be in various forms just as is the case with exothermic reactions.

A few examples of endothermic reactions are:

• Dissolving ammonium nitrate (NH₄NO₃) in water (absorbs heat and cools the surroundings)
• Electrolysis of water to form hydrogen and oxygen gases (absorbs electricity)
• Photosynthesis of chlorophyll plus water plus sunlight to form carbohydrates and oxygen (absorbs light)

 

 

Heats of Reaction

In thermochemistry, we are concerned with the exchange of energy between a chemical system and its surroundings. The form of the energy absorbed or released during a change can vary. It sometimes appears as light, or electrical work, but most often occurs only as heat. When the entire energy change of a reaction involves heat, the amount of heat is called the heat of reaction and is usually represented by the symbol ‘q’.

We show exothermic reactions by q = -ve meaning that energy has been lost from the system. Endothermic reactions are documented by q = +ve meaning that energy was absorbed by the system.

Enthalpy

The first term, enthalpy, refers to the total value of energy of a system when it is at constant pressure. It is symbolized by the letter ‘H’. When a system reacts at constant pressure it will either gain or lose energy and we say that the enthalpy of the system has gone through a change or an enthalpy change, which is symbolized by ΔH.   Δ means “change in”.

ΔH is defined by the equation:        ΔH = H_final – H_inital

H_final is the enthalpy of the system in its final state and H_initial is the enthalpy of the system in its initial state.

For a chemical reaction the above equation can be expressed much more nicely as:

ΔH = H_products – H_reactants

The above equation simply put means “the total heat content of the products minus the total heat content of all the reactants”.

Enthalpy Change and Heat of Reaction at Constant Pressure

When all of the enthalpy change appears as heat, we have a way to measure the enthalpy change for a reaction, because in this circumstance H is equal to the heat of reaction at constant pressure.

ΔH = q (at constant pressure)

The sign of q, defined earlier, is actually determined by the sign of ΔH.   If ΔH is negative so q is negative for exothermic changes. ΔH is positive and so q is positive for endothermic changes.

Simple energy diagrams

A reaction in which heat energy is given off is said to be exothermic.

For an exothermic change:

Notice that in an exothermic change, the products have a lower energy than the reactants. The energy that the system loses is given out as heat. The surroundings warm up

A reaction in which heat energy is absorbed is said to be endothermic.

For an endothermic change:

This time the products have a higher energy than the reactants. The system absorbs this extra energy as heat from the surroundings.

Expressing exothermic and endothermic changes in numbers

Here is an exothermic reaction, showing the amount of heat evolved:

This shows that 394 kJ of heat energy are evolved when equation quantities of carbon and oxygen combine to give carbon dioxide. The mol^-1 (per mole) refers to the whole equation in mole quantities. The reactants (carbon and oxygen) have lost energy during the reaction. When you burn carbon in oxygen, that is the energy which is causing the surroundings to get hotter.

And here is an endothermic change:

In this case, 178 kJ of heat are absorbed when 1 mole of calcium carbonate reacts to give 1 mole of calcium oxide and 1 mole of carbon dioxide. A simple energy diagram for the reaction:

The products have a higher energy than the reactants. Energy has been gained by the system – hence the plus sign.

Measurement of Heat Change

The instrument used for measuring heat change for a chemical reaction is called the calorimeter.

 

Simple Calorimeter

A simple calorimeter for combustion is specifically for determining the heat energy released (given out) for burning fuels. The burner is weighed before and after combustion to get the mass of liquid fuel burned. The thermometer records the temperature rise of the known mass of water. You can use this system to compare the heat output from burning various fuels. The bigger the temperature rise, the more heat energy is released.

Types of Heat Change (Enthalpy reaction)

It is possible to identify heat changes in chemical reactions; it depends on the type of reaction in which it occurs. These heat changes are:

 It is important to note that the thermochemical equation should be balanced in such a way that it represents the formation of one mole of the substance only.

The value of heat of formation at 298 K and 1 atm pressure is called standard heat of formation. Note: The heat of formation of an element in the uncombined state is zero (0).

Example 1: Calculate ΔH for the following reaction:

8 Al(s) + 3 Fe₃O₄(s) –> 4 Al₂O₃(s) + 9 Fe(s)

Heat of Formation Solution:

ΔH for a reaction is equal to the sum of the heats of formation of the product compounds minus the sum of the heats of formation of the reactant compounds:

ΔH = Σ ΔH_f products – Σ ΔH_f reactants

Omitting terms for the elements, the equation becomes:

ΔH = 4 ΔH_f Al₂O₃(s) – 3 ΔH_f Fe₃O₄(s)

The values for ΔH_f may be found in the Heats of Formation of Compounds table

ΔH = 4(-1669.8 kJ) – 3(-1120.9 kJ)

ΔH = -3316.5 kJ

Answer

ΔH = -3316.5 kJ

Example 2: Calculate ΔH for the ionization of hydrogen bromide:

HBr(g) –> H⁺(aq) + Br^–(aq)

Heat of Formation Solution

ΔH for a reaction is equal to the sum of the heats of formation of the product compounds minus the sum of the heats of formation of the reactant compounds:

ΔH = Σ ΔH_f products – Σ ΔH_f reactants

Remember, the heat of formation of H⁺ is zero. The equation becomes:

ΔH = ΔH_f Br^–(aq) – ΔH_f HBr(g)

The values for ΔH_f may be found in the Heats of Formation of Compounds of Ions Table

ΔH = -120.9 kJ – (-36.2 kJ)

ΔH = -84.7 kJ

Answer

ΔH = -84.7 kJ

Heat of combustion (Enthalpy of Combustion)

It may be defined as, “The quantity of heat evolved when one mole of the substance is completely oxidized”.

Eg:

The heat of combustion is very useful in calculating the calorific value of fuels.

 E.g. Combustion of butane evolves -2878.8 k] per mole of heat so it’s calorific value (C.V.) is

Heat of Solution

It may be defined as, “The quantity of heat evolved or absorbed when one mole of a solute is dissolved completely in large excess of water, so that further dilution of solution does not produce any heat change”.

E.g.

Heat of Precipitation

It may be defined as, “The quantity of heat evolved in the precipitation of one mole of a sparingly soluble substance on mixing dilute solutions of suitable electrolytes”.

E.g.

Heat of Vaporisation

It is the change in enthalpy when 1 mole of a substance is converted from liquid to gaseous state at it’s boiling point.

Heat Change and Chemical Bond

Energy changes in chemical reactions result from the breaking and forming of bonds.

The energy needed to break a bond is called the bond-dissociation energy, (DE). In a reaction if the energy needed to break bonds is less than the energy produced by the bonds formed, the reaction will give off energy. Bond dissociation energy, DE, is the amount of energy needed to break a covalent bond

The bond formation energy is the amount of energy released when a covalent bond is formed. These energies are different for each combination of bonds. The weakest bonds have the smallest values. The smaller the bond energy the more reactive and weaker the bond.

Thermodynamics

Thermodynamics is defined as the branch of science that deals with the relationship between heat and other forms of energy, such as work. In other words, thermodynamics is the branch of science which deals with the study of the flow of heat or any other forms of energy into or out of a system as it undergoes a physical or chemical change. It is frequently summarized as three laws that describe restrictions on how different forms of energy can be interconverted.

The Laws of Thermodynamics

First law: Energy is conserved; it can be neither created nor destroyed.

Second law: In an isolated system, natural processes are spontaneous when they lead to an increase in disorder, or entropy.

Third law: The entropy of a perfect crystal is zero when the temperature of the crystal is equal to absolute zero (0 K).

The System and Surroundings

One of the basic assumptions of thermodynamics is the idea that we can arbitrarily divide the universe into a system and its surroundings. The boundary between the system and its surroundings can be as real as the walls of a beaker that separates a solution from the rest of the universe

Internal Energy

One of the thermodynamic properties of a system is its internal energy, E, which is the sum of the kinetic and potential energies of the particles that form the system. The internal energy of a system can be understood by examining the simplest possible system: an ideal gas. Because the particles in an ideal gas do not interact, this system has no potential energy. The internal energy of an ideal gas is therefore the sum of the kinetic energies of the particles in the gas.

The first law of thermodynamics can be captured in the following equation, which states that the energy of the universe is constant. Energy can be transferred from the system to its surroundings, or vice versa, but it can’t be created or destroyed.

First Law of Thermodynamics:

Euniv = Esys + Esurr = 0

A more useful form of the first law describes how energy is conserved. It says that the change in the internal energy of a system is equal to the sum of the heat gained or lost by the system and the work done by or on the system.

First Law of Thermodynamics:

Esys = q + w

The sign convention for the relationship between the internal energy of a system and the heat gained or lost by the system can be understood by thinking about a concrete example, such as a beaker of water on a hot plate. When the hot plate is turned on, the system gains heat from its surroundings. As a result, both the temperature and the internal energy of the system increase, and E is positive. When the hot plate is turned off, the water loses heat to its surroundings as it cools to room temperature, and E is negative.

The relationship between internal energy and work can be understood by considering another concrete example: the tungsten filament inside a light bulb. When work is done on this system by driving an electric current through the tungsten wire, the system becomes hotter and E is therefore positive. (Eventually, the wire becomes hot enough to glow.) Conversely, E is negative when the system does work on its surroundings.

Statement of the first law are:

1. The total energy of an isolated system remains constant though it may change from one form to another.

2. Energy can neither be created nor destroyed, although it can be changed from one form to another.

3. Total energy of a system and surroundings remains constant.

The need for the second law arise from the fact that for a particular process or change, the first law helps us to balance the internal energy, heat released and work done on the system or by the system. But, it does not say anything about the thermodynamic possibility of the process to occur.
The second law explains that ”whenever a spontaneous or irreversible process takes place, it is accompanied by an increase in the total entropy of the universe.” All spontaneous processes take place in the direction of increasing entropy. Entropy is a state quantity that is a measure of the randomness or disorder of the molecules of the system.
The third law does not give any any new concept. It only places a limitation to the value of entropy of a crystalline solid. The entropy of a substance varies directly with temperature. If we increase the temperature of a system, for example water, the molecules attain kinetic energy and starts moving restlessly resulting in an increasing entropy of the system. But, if we cool the system, the vibration of molecules slow down limiting the freedom of movement thereby decreasing the entropy.