1. Introduction to Chemical Energy The Role of Chemical Bonds

2. What is a Chemical Reaction? The process of change of one or more substances into one or more DIFFERENT substances Involves the breaking of EXISITING chemical bonds in starting materials Involves the making of NEW chemical bonds in products

3. What is Chemical Energy? The difference between the energy needed to break bonds and the energy released by making bonds during a chemical reaction Breaking existing bonds always requires energy input Making new bonds always releases energy output Chemical energy is said to be “stored” in chemical bonds and absorbed from or released to surroundings during a chemical reaction

4. Labeling Chemical Energy A process that releases (outputs) energy (usually as heat) is called exothermic A process the absorbs (inputs) energy (usually as heat) is called endothermic Chemical reactions that absorb energy are called endothermic reactions Chemical reactions that release energy are called exothermic reactions

5. Capturing Chemical Energy

6. Breaking and Making Bonds 6CO 2 + 6H 2 O + Energy  C 6 H 12 O 6 + 6O 2

7. Releasing Chemical Energy

8. Breaking and Making Bonds CH 4 + 2O 2  CO 2 + 2H 2 O + Energy

9. Going from reactants to products, energy is either released or absorbed by the chemical system The chemical “system” includes all the matter that is actually involved in the reaction, i.e., undergoing chemical change. Everything else is known as the “surroundings.” This could be the solvent of a solution, the container, or anything in thermal contact with the container.

10. Heat flows either – from the system to the surroundings, or from the surroundings to the system. The chemical “system” includes all the matter that is actually involved in the reaction, i.e., undergoing chemical change. Everything else is known as the “surroundings.” This could be the solvent of a solution, the container, or anything in thermal contact with the container. exothermic endothermic We usually measure the change in temperature of the surroundings to identify which type of energy flow

11. Showing energy changes in a graphical “energy diagram” Reactants  Products Time of reaction Energy in the system (H) Level of energy in reactants Level of energy in products Energy released to surroundings Exothermic Reactions ΔH < 0 Time of reaction Energy in the system (H) Level of energy in reactants Level of energy in products Energy absorbed from surroundings Endothermic Reactions ΔH > 0

12. Calorimetry: How to measure the heat of a chemical reaction System Container (if any) Surroundings (water) Chemical Reaction Calorimeter Insulation

13. Determining Heat Heat cannot be measured directly, but heat (absorbed or released) results in a temperature change, which we can measure. To calculate heat (Q) associated with a temperature change, we use the “heat equation,” Q = m∙C∙∆T where m = mass (in grams), C = specific heat capacity (in J/g- o C or cal/g- o C), and ∆T = final temp – initial temp

14. The “Heat of Reaction” (for chemical reactions) For chemical reactions, the “heat of reaction” (∆H), is calculated as ∆H = heat absorbed or released per mole of a specified reactant, or ∆H = Q / mole The sign of ∆H is negative if the reaction is exothermic – the system has lost heat The sign of ∆H is positive if the reaction is endothermic – the system has gained heat

15. Here’s an example of how to determine the heat of reaction 1.Conduct a chemical reaction in contact with a reservoir of water (calorimeter) -For example: burn a candle under a container of water and measure the temperature change of the water. -Also measure how much wax was combusted by measuring the change in mass of the candle. 2.Calculate Q = (mass of water)(4.18 J/g-oC)(temperature change of water) 3.Here is the chemical equation for the complete combustion of candle wax (C 30 H 62 ): C 30 H 62 + 91/2 O 2  30 CO 2 + 31 H 2 O 4.Calculate n = moles of wax = mass of wax burned / molar mass of wax 5.Calculate ∆H = Q / n and be sure to express with the correct sign

16. Exercise 1P 1. A sample of 0.25 g of propane (C 3 H 8 ) was burned completely in a calorimeter containing 100 mL water. The initial temperature of the water was 25.0 o C and the final temperature was 65.2 o C. What is the heat of combustion, ∆H, for propane in kJ/mol? 2. The complete combustion of carbon is shown by the thermochemical equation: C + O 2  CO 2 + H 2 O, ∆H = - 394 kJ/mol carbon How much total heat (in kJ) would be produced by the complete combustion of 0.5 g of carbon? 3. If the reaction in problem 2 was conducted in a calorimeter containing 100 mL of water at an initial temperature of 22 o C, what was the final temperature of the water in the calorimeter?