Storing Chemical Energy

Transcription

1 Storing Chemical Energy Sugar is nature's way of storing the sun's energy. Light energy is converted into the chemical energy of the bonds in this molecule. Sugar can be converted into many other molecules that we use as fuels. Today we will focus on the amount of energy that is stored in chemical bonds and the ways that the energy can be converted to other forms. Outline Chemical Thermodynamics Using Heats of Formation and Related Quantities Energy in Sugars Homework Chemical Thermodynamics We've discussed aspects of thermodynamics previously. Why cover it again? According to the theoretical physicist Arnold Sommerfied: "Thermodynamics is a funny subject. The first time you go through it, you don't understand it at all. The second time you go through it, you think you understand it, except for one or two small points. The third time you go through it, you know you don't understand it, but by that time you are so used to it, it doesn't bother you anymore." Energy and the First Law of Thermodynamics Energy is conserved. It is neither created nor destroyed. We use this law of thermodynamics all the time. In any chemical reaction, the energy contained in the reactants is equal to the energy contained in the products plus energy released to the environment or absorbed from the environment (as heat or pressure-volume work). In many reactions, energy released or absorbed is entirely heat energy. Under constant pressure conditions, the change in heat energy in a reaction is called the enthalpy change, H. This is easy to measure by using a calorimeter. Once we determine the energy change experimentally from several reactions, we can use the First Law to calculate energy changes in some other reactions. Chemistry 104 Prof. Shapley page 1

2 Entropy and the Second Law of Thermodynamics The entropy of the universe is increasing. Entropy is a measure of randomness or disorder. Spontaneous changes always go from a more ordered state to a less ordered state. There are many examples of this. If you put a drop of blue food coloring in a beaker of water, the dye tends to spread out and you get a pale blue solution. No matter how long you wait, the colored molecules will never come back together to form a drop of dark blue in colorless water. Similarly, all of the oxygen molecules in this room will never spontaneously arrange themselves in one spot, leaving us to be asphyxiated. A block of shiny iron metal exposed to atmosphere will rust but that rust will never spontaneously convert itself back to pure iron and oxygen. It is the entropy of the universe that must increase. This is a combination of the entropy of the system we investigate and the entropy of the surroundings. ΔSuniverse = ΔSsystem + ΔSsurroundings A system can increase order (decrease entropy) but it requires an input of energy from the surroundings, making the surrounding more disordered. For example, you as a living thing are constantly making highly ordered chemical structures within your body. To do this, you require a great deal of energy from your surroundings in the form of food. If you stop taking in this energy, you will ultimately reach a more disordered, decomposed state. Temperature is proportional to the kinetic energy of molecules and atoms. Increasing temperature increases the overall kinetic energy, the random motion of molecules, and so increases entropy. When heat is released by an exothermic chemical reaction to the environment, the entropy of the surroundings increases. This means that ΔSsurroundings will be a positive number. When heat is absorbed from the surroundings by an endothermic chemical reaction, the entropy of the surroundings decreases. This means that ΔSsurroundings will be a negative number. We can calculate the entropy change of the surrounding that is caused by a chemical reaction from the enthalpy change of the reaction and the temperature. ΔSsurroundings = - ΔH/T Chemistry 104 Prof. Shapley page 2

3 Within the chemical reaction system, entropy increases when the number of molecules of products is greater than the number of molecules of reactants or when the products have an inherently greater ability to move (stretch and bend) than the reactants. The change in entropy in a chemical reaction can be calculated based on changes in volume, number of particles, and degrees of freedom between reactants and products. In the forward reaction of the NO2 dimerization reaction, one molecule is formed from 2 molecules and entropy decreases. In the reverse reaction, 2 molecules are formed from 1 molecule and entropy increases. NO2 + NO2 N2O4 Gibbs Free Energy Energy is the capacity to do work or release heat. In chemical reactions, we are not usually concerned with the total energy but rather in changes of the usable chemical energy. This is the Gibbs free energy change or ΔG. The free energy is the capacity to do non-mechanical work at a constant temperature and pressure. Chemical reactions spontaneously proceed in a way that leads to the minimum value of ΔG, that is to the equilibrium condition. When the ΔG of a reaction is a negative number, the reaction will tend to proceed in the forward direction. When the ΔG of a reaction is a positive number, the reaction will tend to proceed in the reverse direction. When the ΔG of a reaction is zero, the system is at equilibrium and there will be no net change. A spontaneous reaction is not necessarily a fast reaction. Remember that the rate of a reaction depends not on the energy difference between product and reactant, but on the height of the activation barrier in a particular pathway. The reaction between molecular oxygen and molecular hydrogen doesn't occur at any measurable rate but it is thermodynamically favorable. At the standard temperature and pressure (STP) of 0 deg C and 1 atmosphere, the Gibbs free energy change is -237 kj/mol. The negative value means that the reaction should be spontaneous in the forward direction. H2(g) + 1/2 O2(g) H2O(l) ΔG 0 = -237 kj/mol At STP the Gibbs free energy is ΔG 0 but the value changes with temperature (T) an pressure (P) according to the formula below where R is the ideal gas constant. ΔG = ΔG 0 + R T ln P Chemistry 104 Prof. Shapley page 3

4 The Gibbs free energy depends on both enthalpy and entropy. The entropy term is small at low temperature but becomes more important as the temperature increases. ΔG = ΔH - T ΔS Equilibrium Chemical equilibrium is a state in which there is no net change in the concentration of reactants and products because the forward and reverse reactions are occurring at the same rate and the Gibbs free energy value is at its minimum. The Gibbs free energy change of a reaction tells us what the concentration of reactants and products will be at equilibrium. For a general reaction of a moles of A, b moles of B and c moles of C combining to form n moles of N, m moles of M, and o moles of O... If the reaction is not at equilibrium, it will proceed in the forward or reverse direction to the minimum value of the Gibbs free energy. At that point, there is no further change in the Gibbs free energy (ΔG = 0) and the reaction is at equilibrium. Using Heats of Formation and Related Quantities Combining Reactions Let's say we are interested in the enthalpy of the reaction that converts 1-butene and molecular hydrogen to butane. Chemistry 104 Prof. Shapley page 4

5 CH 3 CH 2 CH=CH 2 (g) + H 2 (g) CH 3 CH 2 CH 2 CH 3 (g) Because this reaction is generally performed with a catalyst and a high pressure of hydrogen gas, it would be difficult to measure the ΔHrxn under standard conditions by calorimetry. There is a solution! We can more easily measure the heat of combustion of these hydrocarbons. We also know the heat released when hydrogen gas combines with oxygen gas to make water. CH3CH2CH=CH2(g) + 6 O2(g) 4 CO2(g) + 4 H2O(l) CH3CH2CH2CH3(g) O2(g) 4 CO2(g) + 5 H2O(l) H2 + 1/2 O2 H2O(l) ΔHrxn = kj/mol ΔHrxn = kj/mol ΔHrxn = kj/mol If energy is given off in the combustion of butane ( kj/mol), the same amount of energy must be absorbed to make carbon dioxide and water form butane. The means that the reverse reaction will have a ΔHrxn of kj/mol. Eq. 1 CH3CH2CH=CH2(g) + 6 O2(g) 4 CO2(g) + 4 H2O(l) ΔHrxn = kj/mol Eq. 2 4 CO2(g) + 5 H2O(l) CH3CH2CH2CH3(g) O2(g) ΔHrxn = kj/mol Eq. 3 H2 + 1/2 O2 H2O(l) ΔHrxn = kj/mol Add the three reactions above and also add the changes in enthalpy to get the reaction below. CH3CH2CH=CH2(g) + H2(g) CH3CH2CH2CH3(g) ΔHrxn = -147 kj/mol We can combine reactions and their enthalpies to obtain a new reaction and its enthalpy because enthalpy is a state function. It depends on the starting point and the ending point but not on the route traveled. Hess's Law: In going from one set of reactants to one set of products, the enthalpy change is the same whether the reaction takes place in one step or in a series of steps. When we employ this principle, there are two things to keep in mind. (1) Because energy can't be created or destroyed (First Law), if a given quantity of energy is released in one direction, the exact same amount of energy must be absorbed by the reaction in the other direction. (2) When you reverse reactants and products, multiply the ΔHrxn value by -1. The energy units of ΔHrxn are given on a per mole basis. If you multiply the equation by some number, multiply the ΔHrxn value by that number as well. For example: Reaction H2 + 1/2 O2 H2O(l) ΔHrxn (kj/mol) Chemistry 104 Prof. Shapley page 5

6 H 2 O(l) H 2 + 1/2 O H2 + O2 2 H2O(l) Heat of Formation The heat of formation of any compound, ΔHf 0, is the enthalpy of the reaction that forms it from the elements in their most common form at standard temperature and pressure. Data on standard heats of formation of many compounds is widely available and some values are listed here. Let's reconsider the equation for the reaction of 1-butene with hydrogen. CH3CH2CH=CH2(g) + H2(g) CH3CH2CH2CH3(g) 1-butene is formed from graphite, C(s) and hydrogen, H2(g). 4 C(s) + 4 H2(g) CH3CH2CH=CH2(g) ΔHf 0 = 20.4 kj/mol Hydrogen is an element in its common state. H2(g) H2(g) ΔHf 0 = 0.0 kj/mol Butane is also formed from graphite and molecular hydrogen. 4 C(s) + 5 H2(g) CH3CH2CH2CH3(g) ΔHf 0 = kj/mol We could reverse the third equation and add the 3 equations and their ΔHrxn values as we did before. But this is actually the same as adding together the ΔHf 0 values of each product and subtracting the ΔHf 0 values of each reactant from this. Remember that ΔHf 0 values are per mole. If there are 2 (or 3, 4, 5...) moles of a product or reactant in the equation, be sure to multiply the ΔHf 0 value by 2 (or 3, 4, 5...). Entropy of Formation and Free Energy of Formation Entropy and Gibbs free energy are state functions just like enthalpy. We can calculate the entropy change of any reaction if we know ΔSf 0 of all reactants and products, and the free energy change if we know the ΔGf 0 of all reactants and products. Chemistry 104 Prof. Shapley page 6

7 Energy in Sugars Energy from photosynthesis is stored in plants as sugars and polymers of sugars. The chemical composition of this group of molecules, the carbohydrates, is [CH2O]x. All carbohydrate combust to form carbon dioxide and water. These reactions are exothermic with heat energy released from the system. There are MANY carbohydrates with different structures and formulas but, on average, combustion or respiration of a carbohydrate releases 16 kj per gram. The heats of formation of several sugars are listed below. Notice that, in the pictures of the sugar molecules, the carbon atoms and some of the hydrogen atoms are not drawn in. According to convention, a line indicates a bond between 2 saturated carbon atoms unless otherwise specified. We'll discuss this in greater detail next time. Sugar ΔHf 0 (kj/mol) glucose, C6H12O sucrose, table sugar, C12H22O Chemistry 104 Prof. Shapley page 7

8 ribose, C5H10O arabinose, C5H10O Could you use the information in the table above to write a balanced equation for the combustion reactions of these sugars and compute the ΔH rxn? Chemistry 104 Prof. Shapley page 8