Energy. On the ground, the ball has no useful poten:al or kine:c energy. It cannot do anything.

Transcription

1 Energy What is energy? It can take many forms but a good general defini:on is that energy is the capacity to perform work or transfer heat. In other words, the more energy something has, the more things it can do. Most textbooks divide energy into many different types: heat energy, kine:c energy, sound energy, chemical energy etc. In fact, we only need two general categories: kine.c energy and poten.al energy. Kine:c energy is the energy of movement and poten:al energy is stored energy. For instance, sound energy is actually the kine:c energy of oscilla:ng gas par:cles and chemical energy is poten:al energy that arises from the separa:on of electrons and nuclei. Heat is not a special form of energy but is a random mixture of kine:c and poten:al energy. Par:cles are more stable when they have less energy because they have less capacity to change. Ball suspended in the air. Separa:on from the Earth in a gravita:onal field means the ball has poten:al energy. When released, the ball falls and the poten:al energy becomes kine:c energy. On the ground, the ball has no useful poten:al or kine:c energy. It cannot do anything.

2 Thermodynamics Thermodynamics is the study of the interconversion of heat, work and entropy. Thermodynamics tells us a lot about why things happen and what can and cannot happen. It is hard to make a simple defini:on of work as it relates to chemistry. To a very simple approxima:on, we can think of it as being energy that can be directed for a specific ac:on, for example, the energy available to push a piston down. Work is different to heat, as heat is a random mixture of kine:c and poten:al energy and cannot be used directly to perform a specific ac:on. Entropy is a measure of the disorder of a system (more accurately, it refers to the number of possible energy states that can be used to distribute the energy). The concept of entropy can be a liole tricky at first but we shall discuss it in a liole more detail later. There are two forms of thermodynamics commonly used in chemistry. The most common is classical thermodynamics which only works on macroscopic scales for bulk samples of par:cles. The other form is known as sta.s.cal thermodynamics and involves quantum mechanics. We will only look at classical thermodynamics in this course.

3 Thermodynamics Some Ground Rules Classical thermodynamics only applies to a closed system at thermal equilibrium. A closed system does not exchange energy with its surroundings and a system at equilibrium is not hea:ng up or cooling down. a closed system an open system

4 Thermodynamics The First Law The first law of thermodynamics states that the total energy within a closed system will always remain the same. This means that, within the system, we can change energy from one form to another but we cannot lose energy from the total or add energy to the total. O 2 O 2 CO 2 H 2 O CO 2 H 2 O C 8 H 18 C 8 H kj poten:al energy 2715 kj poten:al energy 2715 kj heat energy 5430 kj heat energy

5 Conserva:on of Energy In chemistry, we reword the first law of thermodynamics to say: The total energy of a closed system is the same before, and a<er, a chemical reac.on or phase change. This might seem strange at first. When we burn petrol (gasoline), a lot of heat is produced. However, careful examina:on in a closed environment shows that the heat produced is exactly equal to the poten:al energy lost by the petrol as it turns into water and carbon dioxide.

6 Exothermic and Endothermic Processes When we look at a chemical reac:on or phase change in an open system we can see that some:mes heat comes out of the system and some:mes heat goes into the system. When heat comes out of the reac:on system, we say the change is exothermic. When heat goes into the reac:on system, we say the change is endothermic. Fire is the most well-known exothermic process. Your body also produces heat by a variety of exothermic processes when it oxidises food. Evapora:on is a well known endothermic process. A liquid absorbs energy from from its surroundings when it evaporates, this is how swea:ng cools us down. Instant cool packs for first-aid use an endothermic dissolu.on to quickly cool the temperature of the pack.

7 Enthalpy Changes In chemistry, we usually call heat enthalpy. We use H as the symbol for a change in enthalpy and + H means heat is gained by the system and H means heat is lost by the system. This means that an endothermic process will have a posi:ve H and an exothermic process will have a nega:ve H. The reac:on between hydrogen and oxygen to make water releases a lot of heat. However, the reac:on between carbon and hydrogen to make ethyne (C 2 H 2, also known as acetylene) must absorb heat. The heat lost or gained when making one mole of a compound from its elements is known as the molar enthalpy of forma.on and has the symbol H f. Energy H f ve H 2 + ½O 2 heat out (exothermic) heat in (endothermic) elements H f = 0 product C 2 H 2 2C + H 2 H f +ve H 2 O product

8 A Calorimeter We can measure the enthalpy changes by carefully measuring the quan::es (in moles) of reactants and reac:ng them in an insulated chamber. By carefully measuring the increase or decrease in the temperature of the chamber, we can calculate how much heat has been lost or gained.

9 Hess s Law Whatever path you use to make a certain compound from certain reactants, the overall enthalpy of forma:on will always be the same. This is because the enthalpy of forma:on is a state func.on. Weight is also a state func:on. It does not maoer what order you packed your suitcase in, it will always weigh the same with the same contents and you can calculate the weight of any one thing inside if you know the weight of everything else. This step is very hard to observe but easy to calculate. This means that we can calculate the enthalpy change of a step, even if we cannot measure it directly. We can also calculate theore:cal enthalpy changes for steps that may not even exist in reality.

10 Using Hess s Law If we know the molar enthalpy of forma:on ( H f ) of the reactants and products of a reac:on, we can calculate the molar enthalpy of reac:on ( H rxn ) of that reac:on. For example, take the following reac:on: NO (g) + ½ O 2(g) NO 2(g) We can look up the respec:ve enthalpies of forma:on for NO (g) and NO 2(g). (Remember, elements in their common form at standard temperature and pressure are defined as having a H f of zero). Hess s Law states that the enthalpy of reac:on is equal to the total enthalpy of forma:on of the products minus the enthalpy of forma:on of the reactants. So: H rxn = n products n reactants So if H f NO (g) = kj mol 1 and H f NO 2(g) = 33.2 kj mol 1, ( H f O 2(g) = 0 kj mol 1 ), H rxn = 1(33.2 kj mol 1 ) (1(90.25 kj mol 1 ) + ½ (0 kj mol 1 )) = 57.1 kj mol 1 (3 s.f.) Notes: 1) Mul:ply this value by (no. of moles available/no. moles in reac:on equa:on) to calculate how much energy is released in any given reac:on. For example, if you have 0.5 moles of NO (g), then the energy change is (0.5/1) mol 57.1 kj mol 1 = 28.5 kj. 2) Note the state symbols ( (g), (l), (s) ). Different states have different enthalpies. 3) Watch for limi:ng reagents, e.g. in 2H 2 + O 2 2H 2 O, you might have 5 moles of O 2 and only 5 moles of H 2. In this case, the number of moles of H 2 will limit the total energy released.

11 Bond Strength By combining these calorimetric measurements, we can calculate the energy of individual molecular bonds. We can then see that a lot of energy (measured as heat) is required to break very strong bonds and only a liole energy is required to break weak bonds. Going the other way, forming strong bonds releases a lot of heat and forming weak bonds releases only a liole heat. This explains many things about the stability of substances around us. Many (but not all) reac:ve substances contain at least one weak chemical bond and most (but not all ) unreac:ve substances contain only strong chemical bonds.

12 Phase Changes and Latent Heat When substances change between a solid, a liquid and a gas, they undergo a phase change. During a phase change, the temperature of the substance does not change even though heat is added or released. The heat being added or released at this point is called latent heat. Temperature Heat added

13 Heat Capacity How much heat energy do we have to put in to a sample of a substance to raise its temperature by one kelvin? The amount of heat depends on the molecular nature of the substance and the amount of substance. The molecular nature affects a property known as heat capacity. Specific heat capacity (measured in J g 1 K 1 ) is the amount of heat energy needed to raise one gram of substance by one kelvin. A related quan:ty, the molar heat capacity is the energy required to raise one mole of substance by one kelvin, measured in J mol 1 K 1. Water has a very high specific heat capacity (4.2 J g 1 K 1 ; 75 J mol 1 K 1 ). This means that it takes a lot of heat energy to raise the temperature of water. This is one of the reasons water is so useful for cooling hot processes such as engines. The amount of heat required to convert one mole of liquid to vapour at constant temperature is known as the molar heat of vapourisa.on. It is this energy requirement that cools us down when we sweat. Liquids and gases are useful coolants because they can physically transfer heat by being con:nuously moved from a hot zone to a cooling zone. Water is also useful as a coolant because it is safe, cheap and a liquid over a wide range of temperature.

14 Entropy The loss and gain of energy tells us lot about how and why reac:ons happen in chemistry. However, energy is only half of the picture. The other half is entropy. Entropy is the measure of the dispersal of energy. In other words, the more that energy can disperse itself in a system, the more entropy the system has. The concept of dispersal of energy is open difficult to understand without a deeper understanding the meaning of energy itself, so we shall make a simplifica:on to say that entropy is a measure of disorder, or simply messiness. Simple observa:on of the world around us shows that things get more messy over :me, not more :dy, because non-living things are mindless and do not know what they are doing or where they are going. This tendency towards messiness applies to molecules, and thermodynamics helps us to quan:fy the messiness. The Second Law of Thermodynamics states this behaviour formally: The total entropy in a closed, changing system will tend to increase over :me. In other words, energy will disperse over :me. Another useful way to express the second law is to say: a closed system will always tend towards equilibrium. In other words, energy will disperse un:l it cannot disperse any further, or things will get more disordered un:l they cannot get any more disordered.

15 Entropy: Examples N2H2(l) + 2H202(l) N2(g) + 4H20(g) This reac:on occurs easily (3 molecules become 5 molecules, more messy). 2NH3(g) This reac:on will not occur by itself because entropy decreases too much (4 molecules become 2 molecules, less messy). N2(g) + 3H2(g) However, most reac:ons occur because of a balance between energy and entropy. For instance the reac:on between oxygen and hydrogen to make water causes a loss of entropy (3 molecules become 2 molecules) but breaking three weak bonds to make 4 strong bonds results in a large release of energy. In this case, it is the energy that makes the reac:on happen. O2(g) + 2H2(g) 2H2O(g)

16 What about Life? If entropy must always increase, how can the amazing and ordered structures of life occur by themselves? Remember, the laws of thermodynamics only apply to closed systems. But the Earth is not a closed system, it is an open system and energy con:nuously flows through it from the Sun and from the Earth s own core. When energy flows through an open chemical system, remarkable complexity can arise spontaneously.

17 Reading and Problems Reading Chapter 1: 1-14 Chapter 13: Chapter 15: , (15-6), , Problems Chapter 1: 66 Chapter 13: 44 Chapter 15: 3, 16, 38, 54, 96