Notes: Thermochemistry (text Ch. 16)

Transcription

1 Name Per. Notes: Thermochemistry (text Ch. 16) NOTE: This set of class notes is not complete. We will be filling in information in class. If you are absent, it is your responsibility to get missing information from a fellow classmate or the chemistry website: I. EXOTHERMIC AND ENDOTHERMIC REACTIONS Thermochemistry is the study of heat changes that occur during chemical reactions and physical changes of state. In an exothermic process, heat flows from the system to the surroundings. (Note: the system is the part of the universe on which you are focusing your attention. The surroundings are everything else.) Diagram: In an endothermic process, heat flows from the surroundings to the system. Diagram: Practice: (1) System = person; person is sitting next to a campfire. Is this system exothermic or endothermic? (2) System = person; person is perspiring. Is this system exothermic or endothermic? According to collision theory, atoms, ions, and molecules can react to form products when they collide, provided that the particles have enough kinetic energy. Particles lacking the necessary KE to react bounce apart when they collide, instead of reacting. The minimum amount of energy that particles must have in order to react is called the activation energy. Think of activation energy as a barrier that reactants must cross in order to be converted to products. The activated complex is the name for the arrangement of atoms at the peak of the activation energy barrier. It is an in-between state, being neither reactants nor products. 1

2 In an endothermic reaction, the reactants have less potential energy than the products. H is positive. Energy must be supplied (absorbed by the reactants) to raise the particles up to the higher energy level: In an exothermic reaction, the reactants have more potential energy than the products. H is negative. The extra energy is released to the surroundings: Heat Energy Units The calorie is a commonly used unit for heat energy. Energy contained in the foods we eat is actually measured in kilocalories, or Calories. Example: 1 oreo contains 55 Calories or 55 kilocalories (kcal) calories = 1 Calorie = 1 kilocalorie The Joule (J) is the SI unit for heat energy. We will primarily use Joules in our calculations J = 1 cal 1000 J = 1 kj 2

3 II. ENTHALPY AND THERMOCHEMICAL EQUATIONS The enthalpy (H) of a substance represents the heat content of the material (for systems that occur at constant pressure). The change in enthalpy (H) is the difference in heat content between the products and the reactants of a chemical reaction. Always products reactants!!! Heat Change Sign Convention: Direction of heat flow Heat change Sign of H Reaction Type Out of the system Hproducts - Hreactants < 0 Negative Into the system Hproducts - Hreactants > 0 Positive A thermochemical equation is a chemical equation that gives information about the heat change (H) of the reaction. H is also known as the heat of reaction. Examples of thermochemical equations: (1) 2 H2 (g) + O2 (g) 2 H2O(l) kj Hreaction = 572 kj/mol rxn We know that this reaction is exothermic because the H value appears on the product side (heat is released overall) (2) 2 H2O(l) kj 2 H2 (g) + O2 (g) Hreaction = kj/mol rxn We know that this reaction is endothermic because the H value appears on the reactant side (heat is absorbed overall) Heat of reaction can be calculated by the following formula: Hreaction = Hf products Hf reactants Use the Hf values provided in the notes (at the end) Always products reactants!!! Multiply the Hf value for a compound by its coefficient in the balanced equation Hf for an element (in its standard state) is always zero. Practice: Find the heat of reaction H, and state whether the reaction is exothermic or endothermic. Then write in the H value on the correct side of the equation. (1) CaO (s) + H2O (l) Ca(OH)2 (s) (2) N2(g) + 2 O2 (g) 2 NO2 (g) 3

4 III. SPECIFIC HEAT The specific heat capacity (C), or simply the specific heat, of a substance is the amount of heat it takes to raise the temperature of 1 gram of the substance 1 o C. C = q, where q = heat energy in Joules, m = mass, T = change in temperature m T T = Tfinal - Tinitial Therefore the units for specific heat are J/(g o C) or cal/(g o C). To solve for heat energy, rearrange the equation for q: q = Some substances, such as metals, have low specific heats. This means it doesn t take a lot of energy to cause a temperature change. Other substances, such as water, have high specific heats. It takes more energy to cause a temperature change. On a summer day, why does the concrete deck around a swimming pool become hot, while the water stays much cooler? IV. CALORIMETRY Heat that is released during chemical reactions can be measured by calorimetry. The Law of Conservation of Energy states that energy cannot be created or destroyed, only converted into different forms. We apply LCE when we perform calorimetry experiments. Because of this law, we know that the heat released by a system must be absorbed by the surroundings. Remember: q = heat energy = mct, where m = mass, C = specific heat, and T = change in temp (final initial) Heat released by a system = heat absorbed by its surroundings. mct = mct **You need the negative sign on one side of the equation to make it equal. One side is exo and the other side is endo, so they are opposite in sign. A calorimeter is a device used to measure the absorption or release of heat in chemical and physical processes. For calorimetry calculations, it is assumed that the heat absorbed by the water in the calorimeter equals the heat released by the chemical reaction inside the reaction chamber. It is necessary to measure the mass of the water and the initial and final temperature of the water (q = mct). Diagram of a bomb calorimeter, used to determine heat released during a chemical reaction or heat contained in the foods we eat: (see next page) 4

5 Example: An exothermic reaction between HCl and NaOH was carried out in a calorimeter. The mass of the water in the calorimeter was 150. g. The initial temperature of the water was 25.0 o C. After the reaction was complete, the temperature of the water was 35.0 o C. How much energy in joules and kilojoules did this reaction release? The specific heat of water is 4.18 J/g o C. V. PHASE CHANGES Phase changes are changes in the state of matter of a substance. Phase changes can be exothermic or endothermic processes. Name of phase change States of matter involved? Exothermic or endothermic? Melting Solid to liquid Endothermic Freezing Boiling Condensing Sublimation* Deposition** *Sublimation: solid to gas phase change without passing through the liquid phase (Examples: dry ice, solid air fresheners, mothballs, shrinking ice cubes) **Deposition: gas to solid phase change without passing through the liquid phase (Example: frost on a windshield--water vapor in the air crystallizes on the cold glass) 5

6 Vaporization, evaporation, and boiling: what s the difference? Vaporization is the process by which a liquid changes to a gas. Evaporation and boiling are two types of vaporization. Evaporation is vaporization only at the surface of the liquid, at temperatures below the boiling point. Rate of evaporation depends on temperature, and also on intermolecular forces. A use of evaporation in our bodies is perspiration. How does perspiration help your body cool? How does a fan or a cool breeze help you cool even more? During boiling, vaporization occurs throughout the liquid. The bubbles you see are bubbles of vapor forming from the liquid (it s not air). The pressure inside the bubbles equals atmospheric pressure. The vapor then escapes into the atmosphere. The boiling point of a liquid at a pressure of 1 atmosphere (sea level) is called the normal boiling point. For water, that is 100 o C. The boiling point of a liquid changes as external pressure changes. If the external pressure above the water is higher than normal, the water boils at a higher temperature. If the external pressure is lower than normal, the water boils at a lower temperature. Why don t foods cook the same at high altitude? Example: making spaghetti Phase changes can be represented on a heating curve. The heating curve below is for water. Show where each state of matter exists, label the phase changes, provide values for the temperatures at each phase change and label the direction of the arrows as endo- or exothermic. Assume standard pressure (1 atm). Temperature Tb = o C Tm = o C 6

7 A little more about phase changes: Phase changes always occur at constant temperature. For example, the freezing/melting point of water is 0 o C. If the temperature is exactly 0 o C, there will be a mixture of liquid water and ice present. Because we have both states of matter (solid and liquid) present at the freezing/melting point, we say the solid is in equilibrium with the liquid. If you add heat at this point, you can melt all the ice and then heat the water further if you want. If you take away heat (cool it), you can freeze the rest of the liquid and then cool the ice further if you want. Table of Heats of Formation (continued on next page) Reference: Masterton, Slowinski, Stanitski, Chemical Principles, CBS College Publishing, Note: Heats of formation for elements are always ZERO. Compound ΔHf (kj/mol) Compound ΔHf (kj/mol) AgBr(s) C2H2(g) AgCl(s) C2H4(g) AgI(s) C2H6(g) Ag2O(s) C3H8(g) Ag2S(s) n-c4h10(g) Al2O3(s) n-c5h12(l) BaCl2(s) C2H5OH(l) BaCO3(s) CoO(s) BaO(s) Cr2O3(s) BaSO4(s) CuO(s) CaCl2(s) Cu2O(s) CaCO3(s) CuS(s) CaO(s) CuSO4(s) Ca(OH)2(s) Fe2O3(s) CaSO4(s) Fe3O4(s) CCl4(l) HBr(g) CH4(g) HCl(g) CHCl3(l) HF(g) CH3OH(l) HI(g) CO(g) HNO3(l) CO2(g) H2O(g) H2O(l) NH4Cl(s) H2O2(l) NH4NO3(s) H2S(g) NO(g)

8 H2SO4(l) NO2(g) HgO(s) NiO(s) HgS(s) PbBr2(s) KBr(s) PbCl2(s) KCl(s) PbO(s) KClO3(s) PbO2(s) KF(s) Pb3O4(s) MgCl2(s) PCl3(g) MgCO3(s) PCl5(g) MgO(s) SiO2(s) Mg(OH)2(s) SnCl2(s) MgSO4(s) SnCl4(l) MnO(s) SnO(s) MnO2(s) SnO2(s) NaCl(s) SO2(g) NaF(s) So3(g) NaOH(s) ZnO(s) NH3(g) ZnS(s)