Types of Energy Calorimetry q = mc T Thermochemical Equations Hess s Law Spontaneity, Entropy, Gibb s Free energy

Transcription

1 Unit 7: Energy

2 Outline Types of Energy Calorimetry q = mc T Thermochemical Equations Hess s Law Spontaneity, Entropy, Gibb s Free energy

3 Energy Energy is the ability to do work or produce heat. The energy that is stored in a chemical compound is called chemical potential energy. Heat is energy that flows from a warmer to a colder object. Many chemical reactions involve using or producing heat.

4 Energy When energy flows from a hotter to a colder object, there is a change in temperature. We can measure the energy involved in a chemical reaction using this temperature change. Standard Units of energy (and heat) = Joule, J Common units: calorie, cal; kilocalorie, kcal (Cal) 1 calorie = joules. 1 kcal = 1000 cal

5 Specific heat, c It takes J of energy to raise the temperature of 1 gram of pure water by 1 o C. The amount of energy it takes to raise the temperature of one gram of some substance by one degree Celsius (= 1 Kelvin) is called the specific heat of that substance. (symbol c) Example: It takes less energy to heat up a 100 g piece of concrete than it does to heat up 100 g of water.

6 Specific heat q = m c T Can use this to calculate the heat absorbed or released by a sample. Example: a 5000 g block of concrete increased in temperature by 6.0 o C. The specific heat of concrete is 0.84 J/g. o C. How much heat did it absorb? Textbook p.521 #4-6

7 Calorimetry Calorimetry is the study of the amount of heat released or absorbed during a chemical process. Calorimetry experiments are done using a calorimeter (insulated chamber that has a known mass of water)

8 Calorimetry Eg. 125 g of water in a calorimeter, at 25.6 o C. A metal ball of mass 50.0 g is heated to 115 o C and then dropped into the water. Heat flows from the hot metal to the water until the temperature of the metal and water are equal. The final temperature is measured to be 29.3 o C. What is the heat capacity of the metal? What is the metal? Textbook p.525 #12-15

9 Thermochemistry The study of heat changes that accompany chemical reactions and phase changes is called thermochemistry.

10 Thermochemistry The total energy of a compound is the sum of its chemical potential energy and its kinetic energy, and this amount of energy is called the enthalpy (symbol H), or also the heat content. In chemistry, we are more interested in changes in enthalpy, ΔH.

11 Thermochemistry The change in enthalpy of a reaction (~almost like change in energy of a reaction) is called the enthalpy of reaction/heat of reaction, ΔH rxn. ΔH rxn = H products H reactants

12 Thermochemistry Example: 4Fe(s) + 3O 2 (g) 2Fe 2 O 3 (s) ΔH rxn = kj The product has 1625kJ less energy than the reactants. This much energy has been released to the surroundings.

13 Thermochemistry A thermochemical equation is a balanced chemical equation that includes the physical states of all the species, and also the energy change associated with the reaction, or the change in enthalpy of reaction ΔH rxn.

14 Thermochemistry A reaction that releases heat energy (like an explosion, for example) is called an exothermic reaction. Energy level diagram:

15 Thermochemistry In an exothermic reaction, the reactants have some total energy, and they release it during the reaction, so that the products have a smaller final total energy. So ΔH = negative!

16 Thermochemistry A reaction that absorbs heat energy (like evaporation, for example) is called an endothermic reaction. Energy level diagram:

17 Thermochemistry In an endothermic reaction, the reactants have some total energy, and they absorb some energy from the surroundings during the reaction, so that the products have a greater final total energy. So ΔH = positive!

18 Thermochemistry Reading an energy level diagram: E a = Activation energy. This is the energy needed to break bonds of the reactants, to get the reaction going. So a reaction with a large E a means that a lot of energy is needed to start the reaction.

19 Special Types of Rxns Standard CONDITIONS is 25 o C = 298 K and 1 atm of pressure. (NOT Standard Temperature!) 1. Enthalpy of combustion: energy change when one mole of that substance is burned completely.

20 Special Types of Rxns 2. Enthalpy of formation: energy change when one mole of that substance is formed from its elements. The enthalpy of formation of an element in its standard state (1 atm, 25 o C) is 0

21 Special Types of Rxns 3. Enthalpy of vaporization: ΔH vaporization energy needed to vaporize (liq gas) one mole of a liquid. = - energy released when one mole of gas turns into liquid = - ΔH condensation 4. Enthalpy of fusion: ΔH fusion energy needed to melt one mole of a solid substance. = - energy released when one mole of liquid turns into solid = - ΔH solidification

22 Special Types of Rxns Changing Physical State: 1. Calculate moles of thing changing state. 2. Multiply this by the molar enthalpy of vap/fus/etc. p.532 #23-25

23 Hess s Law You can also calculate the enthalpy change for a reaction if you know the enthalpy change for two or more other reactions that add up to give the reaction that you are trying to solve. Hess s Law: the heat evolved or absorbed in a chemical process is the same, whether the process takes place in one or in several steps.

24 Hess s Law Another way of putting it: Hess s Law: if you can add two or more thermochemical equations to produce a final equation for a reaction, then the sum of the enthalpy changes for the individual reactions is the enthalpy change for the final reaction.

25 Hess s Law Example: Calculate the energy change for the reaction that produces SO 3. 2S(s) + 3O 2 (g) 2SO 3 (g) Given: S(s) + O 2 (g) SO 2 (g) H= -297 kj 2SO 3 (g) 2SO 2 (g) + O 2 (g) H = 198 kj Textbook p. 537 #32-33

26 Bond Enthalpies Bond enthalpy (D) is the amount of energy required to break one mole of a bond. You can also calculate the H rxn if you know the number of bonds broken and formed.

27 Bond Enthalpies When a chemical reaction occurs, bonds in the reactants are broken and new bonds are formed in the products. Bond breaking requires energy. Bond forming releases energy.

28 Bond Enthalpies Bond breaking requires energy. Bond forming releases energy. ΔH rxn = ΣD (Bonds broken) ΣD (Bonds formed)

29 Heating and Cooling Curves Shows the temperature change of a substance as it is heated and changes phase. Temperature increases when there is no phase change. Temperature does not increase during a phase change!

30 Heating and Cooling Curves

31 Spontaneous Processes A spontaneous chemical reaction is one that happens without outside help. Eg. Iron metal rusting: 4Fe + 3O 2 2Fe 2 O 3 2Fe 2 O 3 4Fe + 3O 2 H = kj H = kj

32 Spontaneous Processes So are all exothermic processes spontaneous? Melting of ice is endothermic, and spontaneous! Another factor determines whether a process is spontaneous = entropy.

33 Entropy (S) Entropy, S: the amount of disorder in a system. It is a measure of the freedom of the particles to move, and all the possible arrangements of the particles in a system.

34 Entropy (S)

35 Second Law of Thermodynamics The entropy of the universe always increases for a spontaneous process.

36 Changes in Entropy ( S) 1. Entropy changes can be predicted by thinking about particles in movement. Solid Liquid: S >0 Liquid Gas: S > 0 Solid dissolving into a solvent: Gas dissolving into a solvent:

37 Changes in Entropy ( S) 2. For a chemical reaction, the entropy of a system increases when there are more moles of gaseous product produced than the number of gaseous reactant particles. 2SO 3 (g) 2SO 2 (g) + O 2 (g) 3. Increasing temperature increases the random motion of particles. Textbook p.545 #44-45

38 Gibb s Free Energy ( G) Free energy is the energy available to do work. It combines enthalpy and entropy changes for a process to predict its spontaneity. G = H T S

39 Gibb s Free Energy ( G) The Gibb s Free energy is always negative for a spontaneous process. G = H T S

40 Gibb s Free Energy ( G) The Gibb s Free energy is always negative for a spontaneous process. * Important summary table! Textbook p.548 #46-47