Honors Chemistry 1

1. Honors Chemistry 1 Concept 16 Thermodynamics

2. Chapter 16 Notes Energy and Chemical change.

3. Some background information. • Energy is the ability to do work. Remember the Law of Conservation of Energy? • Chemical Potential Energy (PE) is stored energy.

4. Temperature • Adding or removing energy changes the temperature. • Temperature is the average KE of the atoms and molecules present. • High avg. KE = high temp. • We use relative scales: ºF , ºC and absolute scales: K, R.

5. Heat • Heat is the energy transferred between objects due to the difference in their temperatures. • Heat will naturally go from objects of high KE to objects of low KE. To go in the opposite direction requires work.

6. The many units of heat. • Heat can be measured with many units. • Calorie (cal) is the amount of energy required to raise 1 g of water 1º C. • A kilocalorie (Cal) = 1000 cal. • Joule (J) is the SI unit of heat. 4.184 J = 1 cal.

7. Let’s be specific. • A substance’s specific heat (c) is the amount of energy required to heat 1 g of any substance 1ºC. Every substance absorbs and releases energy at a different rate. This is why 20ºC air is comfortable and 20ºC water isn’t.

8. Calorimetry • To calculate heat absorbed or released use: q= m c  t • at constant pressure, the change in energy of a solution is equal to the mass of the solution times the specific heat capacity of the solution times the change in temperature.

9. Which breaks down into: • Dq = heat absorbed or released (J or cal) • m = mass of substance (g) • c = specific heat of substance (cal or J ) g ºC • Dt = is the change in temperature of the substance (ºC)

10. Equations for Phase Changes. • During heating (temperature change) use Dq = m c Dt • During phase changes use heat of fusion or heat of vaporization. Dq= m DHfus or Dq = m DHvap

11. Heating curve of water • Starting with ice at -4 º C and going to steam at 120 ºC. The total Dq is the sum of the Dq for the individual steps. 120 º C 100 º C steam liquid 0 º C -4 º C ice Adding energy

12. Thermochemistry • The study of the releasing or absorbing of energy during a chemical reaction is thermochemistry. SURROUNDINGS or the ENVIRONMENT HEAT LEAVES SYSTEM EXOTHERMIC SYSTEM HEAT ENTERS SYSTEM ENDOTHERMIC

13. Enthalpy • Enthalpy (H) is the heat content of a system at constant pressure, • You can’t measure enthalpy directly, therefore, you must measure the change in enthalpy during a chemical reactionHrxn = H final - H initial • Enthalpy of reaction: Hrxn = H products- H reactants • For exothermic reactions, Hrxn is negative. • For endothermic reactions, Hrxn is positive.

14. Exothermic Reaction DH < 0 Reactants Surroundings - Hrxn Enthalpy Products

15. Endothermic Reaction DH > 0 Products Surroundings Hrxn Enthalpy Reactants

16. Thermochemical Equations • A balanced chemical equation that includes the physical states of all the reactants and products and also any energy changes is a thermochemical equation. • 4 Fe(s) + 3 O2(g) →2 Fe2O3(s)DH = -1625kJ • Molar ratio still applies 4:3:2:-1625kJ

17. Hess’s Law • Hess's Law -the Enthalpy (ΔHrxn) of a reaction is the sum of the individual ΔH for each step of the reaction mechanism . ΔH for Step #1 + ΔH for Step #2 + ΔH for Step #3 ΔHrxn

18. Spontaneity • Spontaneous reactions occur without any outside intervention. Whether a reaction occurs spontaneously or not is controlled by three factors: • Enthalpy (DH) • Entropy (DS) • Temperature (T)

19. Honors Chemistry 1 Concept 17 Reaction Rates and Equilibriums

20. Chapter 17 Notes Reaction Rates

21. How fast or slow is a reaction? • Let’s look at the reaction • CO(g) + NO2(g)→ CO2(g) + NO (g) • At the start, there are only reactants and no products. As times goes on, there are less reactants and more products. In the end, only product remains. • OK, but how fast did the reaction occur?

22. Reaction Rate • The average rate of a reaction is D quantity =Dmol/L = D concentration D time s D time • We use brackets [ ] to denote concentration in mol/L or M. • Average reaction rate = [NO]t2 - [NO]t1 t2 - t1

23. Collision Theory • In order for a reactants to come together to form products, collisions under the right conditions have to occur. • If the orientation of the collision is correct an activated complex or intermediate will form. • If the collision has sufficient energy (activation energy) then the product is formed.

24. Reaction energy diagram. activated complex activation Energy reactants Energy energy released by reaction products Reaction Progress

25. Factors that affect reaction rate. • Spontaneity isn’t one of them, don’t confuse spontaneous and instantaneous. • The nature of reactants. • Concentration • Surface Area • Temperature • Presence of a catalyst.

26. Nature of the Reactants • Some substances, like halogens and group 1 and 2 elements, are more reactive than others. The more reactive a substance, the faster the reaction occurs.

27. Concentration • In order for a reaction to occur collisions must happen. Obviously, more particles mean more collisions and this means that the reaction occurs faster. • However, too many particles of one kind tend to get in the way of each other and the reaction slows.

28. Surface Area • Increasing the surface area allows the number of collisions between reactants to increase. As with dissolving rates, by grinding, pulverizing, or vaporizing the reactants increases the rate of reaction.

29. Temperature • In general, increasing temperature increases the rate of reaction. The molecules are moving faster and the number of collisions increases. • In general, increasing the temperature by 10 K approximately doubles the rate of reaction.

30. Catalyst or Inhibitor • A catalyst is a substance that increases the rate of reaction without actually taking part in the reaction (it remains unchanged). It does this by lowering the activation energy and thus speeds the formation of intermediate products. • An inhibitor raises the activation energy and slows the reaction down.

31. Reaction Rate Laws • The equation that represents the relationship between the rate and the concentration of reactants is called the rate law. • For the reaction A → B, the rate law is: Rate = k [A], where k is the specific rate constant for that reaction. k is unique for each reaction.

32. Reaction order. • Rate = k [A] is a first order reaction because [A] = [A]1. This means that the concentration and the rate are a direct relationship. • For the reaction a A + b B → products, the general rate law is: Rate = k [A]m [B]n, where m and n are the reaction orders of A and B. Rarely does m=a or n=b. The rate order is m + n.

33. Determining the rate order. • Rate laws are determined experimentally, one method uses initial rates. This is done by comparing the initial rates of a reaction with varying reactant concentration.

34. Instantaneous Reaction Rates. • It sometimes is necessary to know the rate at any given moment, this is the instantaneous rate. • One way is to determine the slope of the tangent to the curve of the reaction at the given time. • Another way is to use the rate law and determine the rate at the given concentration of interest.

35. Reaction Mechanisms • Most reactions consist of a sequence of steps. Each of these is called an elementary step. Several elementary steps make a complex reaction. • The complete sequence of steps in a complex reaction is a reaction mechanism. • The product of the first elementary step of a complex reaction is an intermediate, like catalysts, these don’t appear in the balanced equation.

36. Rate determining step. • In a complex reaction the elementary step that is the slowest determines the rate of the overall reaction. This is the rate determining step. The slowest step always determines the rate.