Energy, Rate, and Equilibrium

1. Energy, Rate, and Equilibrium Dr. Michael P. Gillespie

2. Thermodynamics • Thermodynamics is the study of energy, work, and heat. • We can apply this concept to either chemical or physical changes. • We can calculate the amount of heat obtained from the combustion of one gallon of fuel oil (chemical). • We can calculate the energy consumed or released in the boiling and freezing of water.

3. Kinetic Molecular Theory • Molecules and atoms in a reaction mixture are in constant, random motion. • These molecules and atoms frequently collide with each other. • Only some collisions (those with sufficient energy) will break bonds in molecules. • When reactant bonds are broken, new bonds may be formed and products result.

4. Chemical Reactions and Energy • We cannot measure the absolute value for energy stored in a chemical system. • We can only measure the change in energy as a chemical reaction occurs. • We must establish a boundary between the system and the surroundings.

5. Chemical Reactions and Energy • The system contains the process under study. • The surroundings encompass the rest of the universe. • Energy is lost from the system to the surroundings or energy may be gained by the system at the expense of the surroundings. This usually occurs in the form of heat and can be measured by temperature changes in the system and surroundings.

6. Exothermic and Endothermic Reactions • An exothermic reaction releases energy to the surroundings. The surroundings become warmer. • An endothermic reaction absorbs energy from the surroundings. The surroundings become cooler.

7. Exothermic and Endothermic Reactions

8. Enthalpy • Enthalpy is the term used to represent heat. • The change in enthalpy is the energy difference between the products and reactants of a chemical reaction and is symbolized by ∆H. • Energy released is represented by a negative sign (exothermic) and energy absorbed is represented by a positive sign (endothermic).

9. Enthalpy

10. Spontaneous and Nonspontaneous Reactions • Most, but not all exothermic reactions are spontaneous. • Most, but not all endothermic reactions are not spontaneous.

11. Entropy • The second law of thermodynamics states that the universe spontaneously tends toward increasing disorder or randomness. • Entropy is a measure of the randomness of a chemical system and is represented by the symbol S.

12. Entropy • A random, disordered system has high entropy. • A well-organized system has low entropy. • Disorder or randomness increases when we go from the solid to liquid to gaseous state.

13. Entropy

14. Calorimetry • Calorimetry is the measure of heat energy changes in a chemical reaction. • The device used to measure this is a calorimeter. • The technique involves measuring the change in the temperature of a quantity of water or solution that is in contact with the reaction of interest.

15. Specific Heat • The specific heat of a substance is defined as the number of calories of heat needed to raise the temperature of 1 g of the substance 1 degree Celsius. • Knowing the specific heat of water or another aqueous solution enables the experimenter to calculate the heat released during the reaction.

16. Specific Heat

17. Fuel Value • Fuel value is the amount of energy per gram of food. • The fuel value is generally reported in units of nutritional calories.

18. Nutritional Calorie • One nutritional calorie is equivalent to one kilocalorie (1000 calories). • It is also known as one large Calorie (uppercase C).

19. Bomb Calorimeter • A bomb calorimeter is a special type of calorimeter used to measure the fuel value (Calories) of foods. • It measures the heat released upon combustion of foods.

20. Bomb Calorimeter

21. Kinetics • Thermodynamics help us to decide whether or not a chemical reaction is spontaneous; however, knowing that a reaction can occur spontaneously tells us nothing about the time that it may take. • Chemical kinetics is the study of the rate of the chemical reactions.

22. Kinetics • Kinetic information also tells us something about the mechanism of action. • Kinetic information can provide information about the “shelf life” of processed foods. • The potency of a drug diminishes over time. Kinetics tells us the rate of decline.

23. Chemical Reaction • For a chemical reaction to proceed, sufficient energy must be available to cause the bonds to break. • This energy is provided by the collision of molecules. • Bonds will break and the atoms will recombine in a lower energy arrangement. • An effective collision is one that produces product molecules.

24. Chemical Reaction

25. Activation Energy • Activation energy is the minimum amount of energy required to initiate a chemical reaction. • Enzymes dramatically lower the required activation energy for a chemical reaction to occur.

26. Activation Energy

27. Factors That Affect Reaction Rate • Structure of the reacting species • Concentration of reactants • Temperature of reactants • Physical state of reactants • Presence of a catalyst

28. Structure of the Reacting Species • Bond strengths • Reactions involving ions in solution are usually rapid. Ionic compounds in solution are dissociated, their bonds are already broken. The activation energy should be low. • Reactions involving covalent bonds may proceed more slowly. Covalent bonds must be broken before new bonds are formed. The activation energy is higher.

29. Structure of the Reacting Species • The size and shape of the reactant molecules • Large molecules with bulky groups may block the reactive site.

30. Concentration of Reactants • The rate of a chemical reaction generally increases as the concentration of reactants increases because more reactant molecules results in a greater number of collisions.

31. Temperature of Reactants • The rate of a reaction increases as the temperature increases. • The average kinetic energy of the reacting particles is directly proportional to the Kelvin temperature. • Increasing the speed of the particles increases the likelihood of a collision. • A 10 degree Celsius rise in temperature often doubles the reaction rate.

32. Physical State of Reactants • In the solid state, atoms, ions, and molecules are restricted in their motion. • In the liquid and gaseous states the particles have free motion and close proximity to each other. • Reactions tend to happen fastest in the liquid state and slowest in the solid state.

33. Presence of a Catalyst • A catalyst is a substance that increases the reaction rate. • The catalyst undergoes no net change. • The catalyst creates an alternate chemical pathway for the reactants that requires a lower activation energy.

34. Equilibrium • Equilibrium reactions are ones that do not go to completion. Not all reactants have been converted to products.

35. Equilibrium

36. LeChatelier’s Principle • LeChatelier’s principle states that if a stress is placed upon a system, the system will respond by altering the equilibrium composition in such a way as to minimize the stress.