Chemistry

Key unit Competency:

Explain the concept of energy changes and energy profile diagrams for the exothermic and endothermic processes.

Learning Objectives

By the end of this unit, student will be able to:

• Define the term Thermochemistry.

• Explain the concept of system and distinguish between the types of systems.

• Distinguish between Temperature and heat.

• Explain the concept of Exothermic and endothermic reactions and represent them using energy profile diagrams.

• Carefully deal with reactions that produce a lot of energy.

• Appreciate the use of chemical energy in daily life.

• Exhibit the team working spirit.

• Respect the experimental protocol during chemistry practicals.

• Relate the type of reaction to its energy profile diagram.

• Interprete the experimental results about energy changes occurring during chemical reactions.

• Explain the energy change as a function of the breaking and formation of chemical bonds.

Introductory activity 18

You are provided with the following chemicals and apparatus

• Zinc granules                             

• Dilute hydrochloric acid (HCl)

• Quicklime or Calcium oxide (CaO).

• Ammonium chloride (NH4Cl)                 

• Sodium thiosulphate (Na2S2O3.5H2O)

• Distilled Water

• A thermometer

• 4 Beakers labeled A, B, C and D

• 4 Spatulas

• 4 test tubes

Procedure:

1. In beaker labeled A put dilute hydrochloric acid. In beaker labeled B put water. Read and record the initial temperatures of both solutions.

2. In beaker A put zinc granules using a spatula. In beaker B put calcium oxide using a spatula. Read and record the final temperatures of both solutions.

3. In beaker labeled C put 100 cm3 distilled water. In beaker labeled D put 100 cm3 distilled water Read and record the initial temperatures of both solutions.

4. In beaker C put one endful spatula of ammonium chloride. In beaker D put one endful spatula of sodium thiosulphate. Read and record the final temperatures of both solutions.

Study questions

1.Fill the following table.

          

Which reactions produce heat?

2. Which reactions absorbed heat?

List 2 uses of heat in everyday life.

18.1. Concept of a system

Activity 18.1

Topic: Energy transfer between a system and surroundings.

Apparatus and equipment (per group)

• Eye protection

• Four test-tubes or four expanded polystyrene cups with lids to act as calorimeters

• Spatula

• Teat pipette or small measuring cylinder

• Thermometer

• Access to a balance.

Chemicals (per group)

• Anhydrous copper (II) sulfate. 

• Citric acid crystals (2-hydroxy-1, 2, 3-propane tricarboxylic acid).

• Sodium hydrogencarbonate (baking powder).

• Copper (II) sulfate solution 0.5 mol dm–3.

• Zinc powder.

Procedure

Experiment 1.

1. Put 100 cm³ of water in a test-tube.

2. Record the temperature of the water.

3. Add a spatula measure of anhydrous (white) copper (II) sulfate.

4. Carefully stir, using the thermometer, and record the temperature again.

Safety

• Wear eye protection.

• Anhydrous copper (II) sulfate is harmful.

• Zinc powder is flammable.

Introduction

Instant hot and cold packs are available for use in first aid. This experiment illustrates the types of chemical reaction that occur in these packs.

                   

What to record

What was done and any changes in temperature from the starting temperature of your reaction. A table may be useful.

     

Study questions

1. Identify the reactions that are exothermic and those that are endothermic.

2. Write symbol equations to represent the chemical reaction taking place in Experiment 3.

3. Which two substances could be put in a cold pack?

4. Golfers need a hand warmer to keep their hands warm on a cold day. Which chemicals could be put in these warmers?

All chemical reactions involve the breaking of bonds in the reactants and the formation of new bonds in the products. The breaking of bond requires energy, whereas the formation of bond releases energy.

Thermochemistry is the study of heat and energy associated with a chemical reaction or a physical transformation. Thermodynamics is the study of the relationship between heat, work, and other forms of energy. A reaction may release or absorb energy, and a phase change may do the same, such as in melting and boiling. Energy is exchanged between a closed system and its surroundings during the heating and cooling processes.

A system is a part of the universe which is studied using laws of thermodynamics. Everything outside the system is the surroundings. An infinitely small region separating the system from the surroundings is called boundary. In Chemistry the chemical system consists of reactants and products. The systems are classified according to the number of factors including the composition and the interaction with the surroundings. A system can be homogeneous or heterogeneous. It can be in gaseous, liquid or solid state. A system is said to be in equilibrium when its properties do not change with time. The state of a system is described using its composition, temperature and pressure.

Three types of systems can be distinguished according to the exchange between the system and the surroundings in terms of matter and/ or energy.

1. An open system is a system that can exchange both matter and energy with the surroundings (Figure 18.1).

Examples:    

• All reactions carried out in open containers.

• Evaporation of water in a beaker.

• Hot coffee in a cup.

2. A closed system is a system that can exchange energy but not matter with the surroundings (Figure 18.2).

Examples:

• All the reactions carried out in a closed container. 

• Boiling water in a closed steel vessel. 

• Boiling soup in a closed saucepan.

According to Figures 18.1 and 18.2, both the saucepans without a lid and with a lid, respectively, can absorb heat from the stove and get heated. There is exchange of energy taking place in both cases from the stove (surroundings) to the water (system). However, the saucepan with the lid prevents any change of matter. That is, no matter is added to or removed from the saucepan. On the other hand, in the case of the saucepan without the lid various substances can be added or removed from the saucepan and thus changing the mass of the content. As a conclusion, the lid prevents the exchange of matter between the system and the surroundings. An isolated system is a system which is both sealed and insulated. It can exchange neither matter nor energy with its surroundings.

Examples Hot coffee in a thermos flask (Figure 18.3).The latter is a closed system. The outer surface is insulated and thus neither heat nor matter transfer take place between the system and the surrounding.

Checking up 18.1

1. Which type of thermodynamic system is an ocean? An aquarium? A greenhouse?

2. A closed system contains 2 g of ice. Another 2 g of ice are added to the system. What is the final mass of the system?

3. An isolated system has an initial temperature of 30 oC. It is then placed on top of a Bunsen burner for an hour. What is the final temperature?

4. What type of energy does a pencil on the table have? And what type of energy does a falling pencil have?

5. On which type of system is the first law of thermodynamics is based? What does it stipulates?

18.2. The internal Energy of a system and first Law of Thermodynamics

Activity 18.2

1. a) Calculate the kinetic energy of a running object that has a mass of 80 kg and is running at a speed of 8 m/s.

b) An apple of 154 g is placed in 1.5 m above the ground. Determine its gravitational potential energy? (g = 9.81 m s-2)

c) What is the kinetic energy of a cyclist who, at a certain point in his run down the hill, has a potential energy of 34 300 J and a mechanical energy of 50 725 J?

2. Indicate the direction of heat (from one compartment to another) and explain your answer for the following phenomenon

a) When you touch water in a saucepan on top of a stove with your hand  and you fill it is warm

b) When you touch water from the tap with your hand and you fill it is cold

c) When you mix cold water and warm water

18.2.1. Internal energy

The first Law of Thermodynamics deals with energy that is transferred between a given system and its surroundings in form of heat. The exchange of energy is related to the energy that is stored in the system called internal energy E. The internal energy is the sum of the kinetic and potential energies of the particles that form a system.

 • Kinetic energy (K.E) is the energy possessed by an object in motion such as translation, rotation or vibration.

Where m is the mass of a moving object and v is its speed.                                      

• Potential energy (P.E)

In physics potential energy of an object is defined as energy that an object has because of its position.

Where m = mass in kilograms (kg),   

           g = acceleration of gravity (9.81 m s-2)    

             h = position of the in meters (m). 

In chemistry, the potential energy an object is the energy contained or stored in its chemical bonds.

The total internal energy of a system is the sum of its kinetic energy and its potential energy.

Mathematically, the internal energy (U) of a system is given by the expression:

U = K.E + P.E

18.2.2. Heat energy and temperature

The heat or thermal energy of an object is the total energy of all the molecular motion inside that object. When two bodies are in contact, heat always flows from the object with the higher temperature to that of lower temperature.  Heat transfer ceases when a thermal equilibrium is attained. The heat content of a body will depend on its temperature, its mass, and the material it is made of. Because heat is a form of energy, it is measured in Joules (J) or kilojoules (kJ) or calorie (cal). A calorie is defined as the amount of energy needed to raise the temperature of one gram of water by one degree Celsius.

1 calorie (cal) = 4186 joules (J); 1000 cal = 1 kcal = 4.186 kJ.

The temperature is a measure of the average heat energy (thermal energy) of the molecules in a substance. When an object has a temperature of 100 °C, for example, it does not mean that every single molecule has that exact thermal energy. In any substance, molecules are moving with a range of energies, and interacting with each other. The temperature is a physical measure expressing how an object is hot or cold.The temperature is measured using a variety of temperature scales. The most commonly used are degree Celsius (°C) and degree Kelvin (K):

K = °C + 273

N.B: In thermodynamic calculation, degree Kelvin,  not degree celcius, is used.

First Law of Thermodynamics

Thermodynamics is part of physical chemistry that deals with the relationships between heat and other forms of energy. In particular, it describes how thermal energy is converted to and from other forms of energy and how it affects matter. The first Law of Thermodynamics is a statement about conservation of energy and it categorizes the method of energy transfer into two basic forms: work (W) and heat (Q). The First Law of Thermodynamics states that energy can be converted from one form to another with the interaction of heat, work and internal energy, but it cannot be created or destroyed, under any circumstances. Internal energy refers to all the energies within a given system, including the kinetic energy of molecules and the energy stored in all of the chemical bonds between molecules.

For a closed system (without mass input and output), the internal energy is the sum of the heat energy and the work done by the system or the surroundings

∆U = Q + W

Where W is the energy transferred to the system by doing work and Q is the energy transferred to it by heating.

Let us consider a gas occupying a volume V1 in cylinder with a movable piston on which an external pressure P is applied. If the temperature of the gas increases, it expands and occupies a new volume V2. The change in volume is represented as ∆V, as shown in Figure 18.4. The sign of the work depends on whether it is done by the surroundings on the system or vice versa (Table 18.1).

The work done by the system on the surroundings is negative. Therefore, the first law of Thermodynamics is written as:

ΔU = Q – W

Work (W) is also equal to the negative external pressure on the system multiplied by the change in volume. It can be expressed as:

W = −P∆V

Where P is the external pressure on the system, and ΔV is the change in volume. This is specifically called pressure-volume work. Therefore, the Fist Law of Thermodynamics is expressed using equation: ΔU = Q -P∆V

Table 18.1 Sign convention for Q, W and ∆U

Example:

What is the work of the gas that expands by 10.0 L against an external pressure of exactly 5.5 atmosphere?

Solution

Given that 1atm = 101,325 Pa, 

∆H= -P∆V Pa= 557,287.5 Pa×10.0 L= 5,572,875 J.

N.B: Pascal,

SI unit, is used in calculations.

Glasses P and Q have the same amount of water. Glasses R and S have the same amount of water. The water in Glasses P and R are at the same temperature.

The water in Glasses Q and S are at the same temperature.

1. Fill in the blanks below with the correct answers.

a. The water in Glass……..has the most heat.

b. The water in Glass……..has the least heat.

2. Ari touched a metal spoon. The metal spoon felt cold. Choose the best answer.

a. Heat flows from hand to spoon

b. Heat flows from spoon to hand

c. Heat does not flow

d. Heat flows in both directions

3. Tom placed a metal spoon in a mug of hot coffee as shown below. The metal spoon got hot. Choose the best answer.

a. Heat flows from hand to spoon

b. Heat flows from spoon to hand

c. Heat does not flow

d. Heat flows in both directions

4. Complete the statement below. If two objects are near each other and one object is hotter than the other, then heat will flow from the …………………….object to the………………….. object.

5. Complete the crossword puzzle using the clues given below.

Down

1. Our sense of ………………….cannot measure temperature accurately.

3. Wood is a …………………….conductor of heat.

4. Heat is a form of ………………………..

6. ………………………….is a measure of how hot or cold an object is.

10. Metals can ………………………………when heated.

Across

2. Heat is used to …………………………… food.

5. When two objects of different temperatures are in contact, heat will travel from the ………… object to the other object.

7. The instrument used to measure temperature accurately is a ……………………………..

8. Temperature is measured in the unit ……………………….Celsius (°C).

9. A……………, when used with a temperature sensor, can be used to measure and record temperatures.

10. The Sun is an important ………………………….of heat.

11. A hotter object will has a ……………………….temperature.

12. A gas is compressed and during this process the surroundings does 462 J of work on the gas. At the same time, the gas loses 128 J of energy to the surroundings as heat. What is the change in the internal energy of the gas?

13. What do the first law of thermodynamics have to do with systems?

18.3. Standard Enthalpy changes

Activity 18.3

1. What is meant by standard conditions of temperature and pressure?

2. Which term describes the sum of kinetic energy and potential energy? The standard conditions referring to thermochemical measurements are:

Temperature = 0⁰C or 273 K.

Pressure = 1

atmosphere (atm) or 101, 325 Pa.

The concentration of solutions is 1.0 mol. dm-³ or 1.0 mol L^-1.   

Standard enthalpy change of formation (∆H^o_f)

1. The Standard enthalpy change of formation ∆H^o_f

The standard enthalpy change of formation of a substance is the amount of heat released or absorbed when one mole of that substance is formed from its elements under the standard conditions. It is represented by∆H^o_f

Examples:

Note: The standard enthalpy of formation of substances can be either negative or positive.

The compounds with more negative ∆H^o_f values are more stable than compounds with more positive     ∆H^o_f   values. The standard enthalpy of formation of elements is zero. However, for elements that exist in more than one allotropic form, only the most stable form under standard conditions is given zero standard enthalpy of formation.

2. Standard enthalpy change of Combustion ( ∆H^o_f )

The standard enthalpy change of combustion ( ∆H^oc ) of a substance is the heat evolved when 1 mole of the substance burnt completely in excess oxygen under standard conditions.

Note: For compounds resulting from a direct combination of an element (metal or non-metal) and oxygen, the standard enthalpy of combustion is equal to the standard enthalpy of formation.

3. Standard enthalpy change of neutralization  (∆H^o_n)

The standard enthalpy of neutralization, ΔH^o_n is the enthalpy change which occurs when one gram equivalent of an acid is neutralized by one gram equivalent of a base to produce a salt and water under the standard conditions. The equation of the neutralization reaction is:

4. Lattice enthalpy (ΔH⁰_LE)

The lattice enthalpy is the amount of heat released when one mole of an ionic solid is formed from its gaseous ions under standard conditions. It is also the amount of heat absorbed when an ionic solid dissociates into its gaseous ions under standard conditions.

The lattice enthalpy is negative for the formation of the lattice and positive for the breaking of the lattice. There is a relationship between the size and charge of ions with the lattice enthalpy.

i.The higher the charge of the ions, the higher the strong electrostatic attractions and the higher is the lattice enthalpy.

ii. The smaller is the size of ions, the higher is the strong electrostatic attractions and the higher the lattice enthalpy.

5. Standard enthalpy of Hydration (ΔH^o_Hyd)

The standard enthalpy of Hydration also called Standard enthalpy of solvation is the amount of heat released when one mole of isolated gaseous ions dissolves in

6. The enthalpy change of solution (∆H _s)

The standard enthalpy change of solution is the change in enthalpy that occurs when one mole of a substance is dissolved in water to form an infinitely dilute solution under standard conditions. The dissolution involves the breaking of the lattice into the ions and then there is hydration of these ions. Therefore, the standard enthalpy change of solution is the sum of standard lattice enthalpy and standard enthalpy change of hydration. It may be either positive or negative.

7. Standard enthalpy change of atomization (∆H0_atm)

The Standard enthalpy change of atomization is the amount of heat required to form one mole of free gaseous atoms from its chemical substance under standard conditions. The enthalpy change of atomization is always positive.

Note:

For diatomic molecules, this type of enthalpy is equal to the bond dissociation energy(B.D.E).

Bond energy also called bond enthalpy is a form of potential energy defined as the amount of energy required to break a given chemical bond. It has always positive values which depict the endothermic nature of the bond breaking. The energy required to form a chemical bond is equal in magnitude but opposite in sign to the energy required to break that bond. For example, the energy for breaking a hydrogen-hydrogen bond is 436 kJ.mol-1, and when a hydrogen-hydrogen bond is formed, the process releases 436 kJ.mol-1. In a chemical reaction several bonds are broken and new ones are formed.

• Energy change of reaction = Energy used to break bonds - Energy used to form bonds

Checking up 18.3

1. A gas is compressed and during this process the surroundings does 128 J of work on the gas. At the same time, the gas loses 270 J of energy to the surroundings as heat. What is the change in the internal energy of the gas?

2. Given the following information on mercury, Hg(1 atm), calculate the amount of heat needed at 1 atm to vaporize a 30.0g sample of liquid mercury at its normal boiling point of 357 0C.

Boiling point = 357 °C

Melting point =-38.9 °C

Specific heat (liquid) = 0.139 J/g.°C

∆Hv_ap (357 0C) = 59.3 kJ/mol

∆Hf_us (-38.9 0C)=2.33 kJ/mol

3. Why enthalpy of atomization is always positive?

18.4 Energy profile diagrams for Exothermic and Endothermic reactions.

Activity 18.4

Observe the following image and answer the related questions.

1. Discuss the type of energy form present in points A, B and C of the pathway followed by the vehicle.

2. Discuss how each form of energy changes from point A to point C.

3. Which points corresponds to maximum stability and minimum stability, respectively? Relate your answer to energy concept.

When a chemical reaction happens, the energy is transferred to or from the surroundings and often there is a temperature change. For example, when a bonfire burns, it transfers the heat energy to the surroundings. The objects near the bonfire become warmer and the temperature rise can be measured with a thermometer.

There are some chemical reactions that must absorb energy in order to proceed. These are endothermic reactions. Some other chemical reactions release energy to the surroundings. The energy released can take the form of heat, light, or sound. These are exothermic reactions.

1. Exothermic reactions

They are characterized by an increase in the temperature of the surroundings, i.e. energy is given up. Heat is lost to the surroundings and by convention it is negative and represented as: ΔH < 0

For exothermic reaction (Figure 18.8), total energy of the reactants is higher than in the product, because the heat energy absorbed during bond breaking is lower than the heat energy released during bond formation.

Examples of exothermic reactions are:

1.Burning different substances

2. Neutralization reactions between acids and alkalis

3. The reaction between water and calcium oxide

4. Termite reaction: This is the reduction of metal oxides in which a large amount of heat is liberated. It is very useful for the connecting of broken metal parts. When Aluminum powder reacts with iron oxide or chromium oxide, a large amount of heat is released (about 3500 °C is attained to weld broken metallic parts.

5. The reaction of sodium and chlorine to yield table salt is an exothermic reaction. This reaction produces 411 kJ of energy for each mole of salt that is produced:

2. Endothermic reactions

These are reactions that take place by absorbing the energy from the surroundings. The energy is usually transferred as heat energy; in this case the surroundings loses energy to the reactants causing the surroundings to get colder. Endothermic reactions cannot occur spontaneously. Work must be done in order to get these reactions to occur. When endothermic reactions absorb energy, a temperature drop in the surroundings is observed during the reaction. Endothermic reactions are characterized by positive heat flow (into the reaction) and an increase in enthalpy, by convention it is represented by: ΔH > 0

For endothermic reaction (Figure 18.9), the total energy of the reactants is lower than the product, because the heat energy absorbed during bond breaking is higher than the heat energy released during bond formation.

You have certainly experienced this effect when you put a drop of methanol or any other volatile substance on your skin; you feel cold because that part of your skin is supplying energy to evaporate the volatile liquid.

Examples

1. Water evaporation  

2. The thermal decomposition of calcium carbonate to produce Quicklime, CaO.

3. Cooking

3. Activation energy, E_a

The activation energy is the minimum energy required for a chemical reaction to take place. It is the energy barrier that has to be overcome for a reaction to proceed. Without that minimum energy, the reaction will not take place. That is why, for example, the only fact that a dry wood is in contact with oxygen of air will not start burning; there is a need of supplying the minimum energy to overcome the activation energy barrier, this is done by using a burning match.

4. Activated complex

The activated complex is the intermedicate species, where former chemical bonds are being broken, whereas new chemical bonds are being formed. In term of energy, it corresponds to the activation energy.

Examples

1. Let us consider the reaction between hydrogen and fluorine to form hydrogen fluoride.

Determine the enthalpy change of the reaction and decide whether the reaction is endothermic or exothermic.

Data: The bond energies of H-H, F-F, and H-F are 436 kJmol-1,155kJmol-1 and 567 kJmol-1, respectively.

Solution:

The stoichiometric coefficients show that a hydrogen-hydrogen bond and a fluorinefluorine bond are broken. Moreover, two hydrogen-fluorine bonds are formed. The overall energy change for this process is tabulated below.

2. Consider the complete combustion of butane and answer to the related questions. Given the bond energies of reactants and products in the following table.

a. Use bond energies to estimate the enthalpy change for the said reaction.

b. What is the nature of the reaction? Explain.

Solution:

a. The balanced equation for the reaction is:

Referring to the values of bond energy in the Table 18.1 and taking into consideration  the stoichiometric coefficients, the total bond enthalpy calculated is given as follows:

Checking up 18.4

Use the following  potential energy diagram to answer the questions below:

1.What is meant by Activation energy?

2. Determine the energy of the reactants.

3. Determine the energy of the products.

4. Determine the activation energy for the forward reaction.

5. Determine the activation energy for the reverse reaction.

6. Determine the enthalpy change of reaction for the forward reaction.

7. Determine the enthalpy change of reaction for the reverse reaction.

8. Fill in using exothermic or endothermic.

a. The forward reaction is ……………………..

b. The reverse reaction is ………………………

9. Which chemical species or set of chemical species represent the activated complex?

10. Which one of the chemical bonds A-X and M-X is stronger? Explain.

11. State the chemical species whose particles move the fastest. Explain your answer.

12. State the chemical species whose particles move the slowest. Explain your answer.

13. The compound AX and the element M are in gaseous and solid states, respectively. What effect would grinding M into a fine powder have o the above graph?

18.5. End unit Assessment

Observe the diagram hereafter answer the related questions

. Regarding the absorption or release of energy, what is the nature of the overall
reaction?
b. What is the activation energy for the forward reaction?
c. What is the activation energy for the reverse reaction?
d. Determine the enthalpy change of reaction for the forward reaction?

e. Is the reverse reaction endothermic or exothermic?
f. Which chemical species constitute the activated complex?
g. Which chemical species or set of chemical species have the maximum potential
energy?
h. Which chemical species or set of chemical species have the maximum kinetic energy?
i. Which chemical species or set of chemical species have the strongest bonds?
j. Which chemical species or set of chemical species have the weakest bonds? k. What is the enthalpy change of reaction for the reaction X2Y2 X2 + Y2 ? As reactant particles approach each other before collision the potential energy goes whereas the kinetic energy goes ……………………........
l. Which one of the forward reaction and reverse reaction is more likely to be faster?
m. As particles of the newly formed product move away from one another their potential energy goes …………………….......while their kinetic energy goes
……………………..
n. State the meaning of the term Activated complex.
o. Which chemical species or set of chemical species correspond
P.  What is the effect of caltalyst if it is added to the above reaction?

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