Topic 5 - Thermodynamics (Part 3)-1.pptx

Thermodynamics

1. General chemistry 2 Topic 5 – Thermodynamics (Part 3)

2. Objectives Students should be able to: Define the term enthalpy and enthalpy change Interpret energy profile diagrams for exothermic and endothermic reactions Describe the process for measuring enthalpy changes for reactions Describe the conditions for measuring standard enthalpy changes Define standard enthalpy changes of neutralization and solution Calculate standard enthalpy changes of neutralization and solution

3. What is enthalpy?

4. Enthalpy Since most chemical reactions happen in open systems, chemists ordinarily use a property known as  enthalpy ( H )  to describe the thermodynamics of chemical and physical processes. Enthalpy is defined as the sum of a system’s internal energy ( U ) and the mathematical product of its pressure ( P ) and volume ( V ). Enthalpy is also a state function. Enthalpy values for specific substances cannot be measured directly; only enthalpy changes for chemical or physical processes can be determined. ΔH=ΔU+PΔV For processes that take place at constant pressure (a common condition for many chemical and physical changes), the enthalpy change (ΔH) simplifies to: ∆ H sys = q p (where q p is the heat of reaction under conditions of constant pressure) And so, if a chemical or physical process is carried out at constant pressure with the only work done caused by expansion or contraction, then the heat flow (q p ) and enthalpy change (ΔH) for the process are equal.

5. Types of chemical reactions and Enthalpy sign Exothermic reactions – the system releases energy to the surroundings and the temperature of the surroundings increases as a result e.g combustion, neutralization reactions A negative value of an enthalpy change, Δ H  < 0, indicates an exothermic reaction Endothermic reactions – the system absorbs energy from the surroundings and the temperature of the surroundings decreases as a result e.g thermal decomposition reactions, dissolution of some salts A positive value, Δ H  > 0, indicates an endothermic reaction. Bond making is an exothermic process and as such its energy values are negated (- ve ) Bond breaking is an endothermic process and as such its energy values are positive (+ ve )

6. Exothermic Energy/Enthalpy Profile Diagrams Key features: Relative energy and formulae of reactants and products on y-axis Reaction pathway on x-axis Energy of products is LESS than energy of reactants Enthalpy change is – ve Energy is lost to the surroundings in the form of heat Down arrow indicates energy released

7. Endothermic Energy/Enthalpy Profile Diagrams Key features: Relative energy and formulae of reactants and products on y-axis Reaction pathway on x-axis Energy of products is GREATER than energy of reactants Enthalpy change is + ve Energy is absorbed from the surroundings in the form of heat Up arrow indicates energy absorbed

8. Examples of Enthalpy Diagrams Enthalpy of Combustion of Methane Enthalpy of Solution of Ammonium Nitrate

9. Calculating Enthalpy Changes When 100 mL of 0.200 M NaCl( aq ) and 100 mL of 0.200 M AgNO 3 ( aq ), both at 21.9°C, are mixed in a coffee cup calorimeter, the temperature increases to 23.5 °C as solid AgCl forms. How much heat is produced by this precipitation reaction? What assumptions did you make to determine your value?

10. Standard Enthalpy Changes Standard enthalpy changes refer to reactions done under standard conditions , and with all reactants and products present in their standard states . Standard states are sometimes referred to as "reference states". The symbol for a standard enthalpy change is ΔH ϴ

11. Standard conditions and states Standard conditions are: 298 K (25°C) a pressure of 1 bar (100 kPa ). where solutions are involved, a concentration of 1 mol dm -3 Standard states This is the physical and chemical state that you would expect to find reactants in under standard conditions. That means that the standard state for water, for example, is liquid water, H 2 O(l) -  not  steam or water vapour or ice. Oxygen's standard state is the gas, O 2 (g) -  not  liquid oxygen or oxygen atoms.

12. Enthalpy changes that can be determined by experimentation The following enthalpy changes can be measured using a coffee-cup/simple calorimeter: Standard enthalpy change of reaction – can be exothermic or endothermic Standard enthalpy change of neutralization – always exothermic Standard enthalpy change of solution – can be exothermic or endothermic A bomb calorimeter is used to measure: Standard enthalpy change of combustion – always exothermic

13. Measuring standard enthalpy changes for reactions using calorimetry

14. General procedure for measuring enthalpy changes using a calorimeter A known amount of reactant is added to the calorimeter. The change in temperature of the water/solution in the calorimeter is recorded using a thermometer. The energy transferred is then calculated. The enthalpy change is then calculated.

15. General assumptions made when calculating ΔH° from experimental data Since most of the reactions will involve solutions, we can proceed as if it were water in terms of its specific heat and mass values. The specific heat of water is approximately 4.184 J/g °C, so we use that for the specific heat of the solution. Water has a density of 1g/cm 3 , so 1g = 1 cm 3 We also assume that the only heat transfer that occurs is between the system and the solution and any transfer to the calorimeter itself is negligible.

16. Standard enthalpy change of neutralization, ΔH° neu Definition - The enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water . Enthalpy changes of neutralization are always negative. Equation: Net ionic equation?

17. Determining ΔH° neu experimentally

18. Calculating ΔH° neu Equation: D H = - (mc Δ T/n) Measurements required: D H = energy transferred from reaction(system) to solution (surroundings) m = total mass of solutions in the calorimeter in grams (assume solution has the same density as water, 1 ml = 1 g) Δ T = change in temperature c = specific heat capacity of the solution (assume solution has the same specific heat capacity as water = 4.184Jg -1 K -1 ) n = # of moles of H 2 O produced

19. Worked Example #1 Calculate the enthalpy change of neutralization for the reaction between 50ml of 1M HCl and 50ml of 1M NaOH if a temperature increase of 6.2 o C is observed. The value given for this reaction from a data book was -57.1 kJ/mol, suggest a reason for the difference between your calculated value and that in the data book.

20. Practice Question When 150 cm 3  of 1M dilute nitric acid is neutralized by 150 cm 3  of 1M sodium hydroxide, there is a temperature rise of 6.81°C. Calculate the standard enthalpy of  neutralisation  of nitric acid.

21. Practice Question #2 Assume you mix 200 mL of 0.400M HCl with 200 mL of 0.400M NaOH in a coffee-cup calorimeter. The temperature of the solutions before mixing was 25.10 ˚C. The standard enthalpy change of neutralization is -56.01kJ/mol. Calculate the final temperature of the solution after mixing.

22. Worked Example #2 50.0 cm 3 of 1.50 mol dm –3 sodium hydroxide is mixed with 100.0 cm 3 of 1.00 mol dm –3 hydrochloric acid. Both solutions were initially at 19.3  C and when they were mixed the temperature rose to a maximum of 25.3  C.

23. Practice Question If 50.00 mL of 1.05 M NaOH is added to 25.00 mL of 1.86 M HCl, and the temperature rose from 24.72°C to 32.45 o C, what is the standard enthalpy change of neutralization for the reaction?

24. Practice Question If 50.00 mL of 1.05 M NaOH is added to 25.00 mL of 1.86 M HCl, and the temperature rose from 24.72°C to 32.45 o C, what is the standard enthalpy change of neutralization for the reaction?

25. Standard enthalpy change of solution, ΔH° sol Definition - The enthalpy change when 1 mole of a substance dissolves in water to give a solution of infinite dilution or 1 M under standard conditions. Example: NH 4 NO 3   (s)    + aq →   NH 4 +   (aq)  + NO 3 -   (aq)

26. Calculating ΔH° sol Equation: D H = (mc Δ T/n) Choose appropriate sign based on information given. Measurements required: D H = energy transferred from reaction(system) to solution (surroundings) m = mass of water in the calorimeter in grams (1 ml = 1 g) Δ T = change in temperature c = specific heat capacity of water = 4.184Jg -1 K -1 n = # of moles of solid dissolved

27. Worked Example #3 Calculate the enthalpy of solution of NaOH when 6.2g of NaOH dissolves in 100ml of water and gives an increase in temperature of 16 o C. Mr of NaOH = 40 g/ mol

28. Practice Question When solid ammonium nitrate dissolves in water, the solution becomes cold. This is the basis for an “instant ice pack” . When 3.21 g of solid NH 4 NO 3  dissolves in 50.0 g of water at 24.9 °C in a calorimeter, the temperature decreases to 20.3 °C. Determine the standard enthalpy of solution of the ammonium nitrate. Mr of ammonium nitrate = 80 g/ mol

29. Practice Question When solid ammonium nitrate dissolves in water, the solution becomes cold. This is the basis for an “instant ice pack” . When 3.21 g of solid NH 4 NO 3  dissolves in 50.0 g of water at 24.9 °C in a calorimeter, the temperature decreases to 20.3 °C. Determine the standard enthalpy of solution of the ammonium nitrate. Mr of ammonium nitrate = 80 g/ mol