THERMODYNAMICS 11/14/2018 THERMOCHEMISTRY OBJECT OF THE THERMODYNAMICS FUNDAMMENTAL ASPECT OF THERMODYNAMICS

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1 THERMODYNAMICS Thermochemistry In thermodynamics we study the energy changes that accompany physical and chemical processes. Usually these energy changes involve heat hence the thermo- part of the term. There are the two main aspects of thermodynamics. OBJECT OF THE THERMODYNAMICS 2 The first aspect is thermochemistry. This practical subject is concerned with how we observe, measure, and predict energy changes for both physical changes and chemical reactions. THERMOCHEMISTRY The second aspect is addressed to a more fundamental aspect of thermodynamics. How to use energy changes to tell us whether or not a given process can occur under specified conditions to give predominantly products (or reactants) how to make a process more (or less) favourable. FUNDAMMENTAL ASPECT OF THERMODYNAMICS 3 4 1

2 Thermochemistry deals with changes in heat during chemical reactions. A main goal of the thermochemistry study is to determine the quantity of heat exchanged between a system and its surroundings. The system is the part of the universe being studied, while the surroundings are the rest of the universe that interacts with the system. Definitions System and surroundings 5 6 An open system is a system that freely exchanges energy and matter with its surroundings. System and surroundings Open system 7 8 2

3 Open systems A closed system exchanges energy but not matter with an outside system (surroundings). Although it is typically portion of larger system, it is not in complete contact with it. Closed system 9 10 An isolated system can exchange neither energy nor matter with its surroundings (an outside system). While it may be portion of larger system, it does not communicate with the outside in any way. Examples of such system type are: physical universe and a closed thermos bottle (though its isolation is not perfect). Isolated system

4 13 Comparison of systems 14 Q, A closed system contains 2g of ice. Another 2g of ice are added to the system. What is the final mass of the system? Q An isolated system has an initial temperature of 30 o C. It is then placed on top of a bunsen burner for an hour. What is the final temperature? Q. Identify system type (open, closed or isolated) from description below and fill the empty space in the table below System description System type Coffee in perfectly closed Thermos flask Combustion of gasoline in car engine Mercury in thermometer Living plant Quiz Electric battery

5 Q. Which type of thermodynamic system is: 1. an ocean? 2. an aquarium? 3. a pizza delivery bag? 4. a greenhouse? 5. a man? Matter and ENERGY energy Matter is anything that has a mass and occupies some space. All bodies consist of a matter. Mass is a measure of the quantity of a matter in a sample of any material. The more massive an object is, the more force is required to put it in a motion. Energy is measure of the ability of a body or system to do work or produce any change. No activity is possible without energy. MATTER AND ENERGY Energy

6 Energy can take many forms: electrical energy, radiant energy (light), nuclear energy, chemical energy. ENERGY FORMS Commonly we classify energy into two general types: kinetic and potential. Kinetic energy is the energy of motion. The kinetic energy of an object is equal to one half its mass, m, times the square of its velocity, v. 1 2 E k m v 2 The heavier a hammer is and the more rapidly it moves, the greater its kinetic energy and the more work it can accomplish. KINETIC ENERGY Q. The kinetics energy of solid body with the mass of 5 kg which moved with speed 8 m s -1 is equal: a) 40 kg m s -1 ; b) 320 J; c) 160 J d) 160 kg m 2 s -2 Potential energy (E P ) is the energy that a system possesses by virtue of its position or composition. It is stored energy. The work that we do to lift an object is stored in the object as energy. E p = m g h Where: m mass (kg); h body movement (change of height, m); g gravitational acceleration, (10 m s -2 ) Kinetics energy in questions POTENTIAL ENERGY

7 Q. When a bucket with 10 kg of water is picked up at the height of 1 m the potential energy is as follows: a) 10 kg m; b) 100 J; c) 100 kg m 2 s -2 ; d) 100 kg If we drop a hammer, its potential energy is converted into kinetic energy as it falls, and it could do work on something it hits for example, drive a nail or break a piece of glass. Potential energy in question EXAMPLE: Ep Ek Q. What type of energy does a stationary pencil contain? falling pencil? Chemical changes always involve energy changes. However, some energy transformations do not involve chemical changes at all. For example, heat energy may be converted into electrical energy or into mechanical energy without any simultaneous chemical changes

8 With the dawn of the nuclear age in the 1940s, scientists, and then the world, became aware that matter can be converted into energy. In nuclear reactions, matter is transformed into energy. The relationship between matter and energy is given by Albert Einstein s now famous equation: E = m c 2 The Law of Conservation of Matter and Energy Many experiments have demonstrated that all of the energy involved in any chemical or physical change appears in some form after the change. These observations are summarized in the Law of Conservation of Energy: Energy cannot be created or destroyed in a chemical reaction or in a physical change. It can only be converted from one form of energy to another. Law of Conservation of Energy Now that the equivalence of matter and energy is recognized, the Law of Conservation of Matter and Energy can be stated in a single sentence: The combined amount of matter and energy in the universe is fixed. The Law of Conservation of Matter and Energy

9 Energy is very important in every aspect of our daily lives. The food which we eat supplies the energy to sustain life with all of its activities and concerns. The availability of relatively inexpensive energy is an important factor in our technological society. MEANING OF THE ENERGY IN THE LIFE The concept of energy is at every heart of science. All physical and chemical processes are accompanied by the transfer of energy. Energy cannot be created or destroyed. We must understand how to do the accounting of energy transfers from one body or one substance to another or from one form of energy to another. ACCOUNTING OF ENERGY We can define energy as follows: Energy is the capacity to do work or to transfer heat. The amount of heat transferred in a process is usually expressed in joules or in calories. The SI unit of energy and work is the joule (J), which is defined as 1 kg m 2 /s 2. Not SI unit of energy e.g. 1 cal = J ENERGY - DEFINITION Heat, energy, work units

10 Heat is a form of energy that always flows spontaneously from a hotter body to a colder body never in the reverse direction Heat transfer concerns the generation, use, conversion, and exchange of heat (thermal energy) between physical systems. Heat transfer is classified into various mechanisms, such as thermal conduction (diffusion), thermal convection, thermal radiation and by phase changes. Heat as a form of energy 38 Heat transfer 39 Temperature measures the intensity of a heat, the hotness or coldness of a body. A piece of metal at 100 C feels hot to the touch, whereas an ice cube at 0 C feels cold. Why? Because the temperature of the metal is higher, and that of the ice cube lower, than body temperature. Conduction is the transfer of thermal energy through direct contact between particles of a substance, without moving the particles to a new location. Convection is the transfer of thermal energy through movement of particles from one location to another Radiation is the emission of energy as waves or particles or rays. TEMPERATURE vs. HEAT

11 State functions It turns out that the energy of an object depends only on the object s current condition. The complete list of properties that specify an object s current condition is known as the state of the object. In chemistry it is usually enough to specify the object s pressure, temperature, volume, and chemical compositions (numbers of moles) to give the state of the object Any property of a system that depends only on the values of its state functions is also a state function. For instance, the volume of a given sample of water depends only on temperature, pressure, and physical state; volume is a state function. The thermodynamic state of a system is defined by a set of conditions that completely specifies all the properties of the system. We shall encounter other thermodynamic state functions. STATE FUNCTION THERMODYNAMIC STATE

12 This set commonly includes: temperature, T, pressure, P, volume, V, composition (identity and number of Ice liquid water moles of each component), n, physical state (gas, liquid, or solid) of Changes in physical state of water due to temperature changes each part of the system. THERMODYNAMIC STATE Steam The properties of a system such as P, V, T are called state functions. The value of a state function depends only on the state of the system and not on the way in which the system came to be in that state. A change in a state function describes a difference between the two states. 3 Indirect 1 Indirect Direct STATE FUNCTIONS

13 For instance, consider a sample of one mole of pure liquid water at 30 C and 1 atm of pressure. If at some later time the temperature of the sample is 22 C at the same pressure, then it is in a different thermodynamic state. Thus change in temperature is equal to: t = t final - t initial We can tell that the net temperature change is 8 C. It does not matter whether: (1) the cooling took place directly (either slowly or rapidly) from 30 C to 22 C, or (2) the sample was first heated to 36 C, then cooled to 10 C, and finally warmed to 30 C, then cooled to 22 o C or (3) any other conceivable path was followed from the initial state to the final state. STATE FUNCTIONS STATE FUNCTION The change in other properties (e.g., the pressure) of the sample is likewise independent of path. The most important use of state functions in thermodynamics is to describe changes. We describe the difference in any quantity, X, as X X FINAL X INITIAL When X increases, the final value is greater than the initial value, so ΔX is positive; a decrease in X makes ΔX a negative value. STATE FUNCTION STATE FUNCTIONS

14 We can determine the energy change associated with a chemical or physical process by using an experimental technique called calorimetry. This technique is based on observing the temperature change when a system absorbs or releases energy in the form of heat. This is in turn the effect of chemical or physical process under study. Thermochemistry CALORIMETRY, thermochemistry The experiment is carried out in a device called a calorimeter, in which we measure the temperature change of a known amount of substance (often water) which specific heat is known. The specific heat, c, is the amount of heat per unit of mass required to raise the temperature by one degree Celsius or Kelvin with no change in phase. Specific heat

15 SPECIFIC HEAT [c] The specific heat of each substance, a physical property, is different for the solid, liquid, and gaseous phases of the substance. For example, the specific heat of: ice is 2.09 J/g C near 0 C; liquid water is 4.18 J/g C; steam is 2.03 J/g C near 100 C. SPECIFIC HEAT Specific heats and molar heat capacities for various substances at 293 K (20 o C) Substance c in J/g K Molar C J/mol K Aluminum Copper Gold Lead Silver Zinc Mercury Alcohol(ethyl) Water Ice (-10 C) SPECIFIC HEAT 61 Q. How many heat is needed to heat up 10 g of liquid water from 10 o C to 40 o C. [c of liquid water is 4.18 J g -1 o C -1 ]. a) 418 J; b) 1254 J; c) 1672 J 62 15

16 Answer: Q = m c ΔT Q = (10 g). (4.18 J g -1 o C -1 ) (40 10) o C The specific heat, c, is the amount of heat per unit of mass required to raise the temperature by one degree Celsius or Kelvin with no change in phase. Q = = 1254 J Specific heat Changes in phase (physical state) absorb or liberate relatively large amounts of energy (see Figure NEXT SLIDE). Changes in phase

17 A hot object, such as a heated piece of metal (a), is placed into cooler water. We say that they are then at thermal equilibrium Heat is transferred from the hotter metal bar to the cooler water until the two reach the same temperature (b). Quiz A 34 gram piece of an unknown metal absorbs Joules of energy when the temperature increased from 10 o C to 32 o C. What is the specific heat of the substance? Hint: You are solving for Specific Heat (Cp) not heat absorbed. THERMAL EQUILIBRIUM Q. A 385 grams chunk of iron is heated to 97.5 o C. Then it is immersed in 247 gram of water originally at 20.7 o C. When thermal equilibrium has been reached, the water and iron are both at 31.6 o C. Calculate the specific heat of iron. Solution: The number of heat gained by water from temperature 20.7 o C to 31.6 o C =the amount of heat which is lost by the iron. Q water = (247 g). (4.18 J g -1 o C -1 ) ( o C) Q water = (247). (4.18). (10.9) = J

18 If a 25.2 g piece of silver absorbs 365 J of heat, what will be the final temperature of the silver if the initial temperature is 22.2 o C? The specific heat of silver is J/g K Expansion work. Pressure-volume work of gas Atmospheric pressure Atmospheric pressure Run Pin position Initial bucket temperature ( o C) Final bucket temperature ( o C) 1 locked pin 2 unlocked Gas under Gas at atmospheric pressure pressure Constant pressure 1 2 Heat capacity of calorimeter = kj/ o C, mesured in separated experiment

19 q for run 1 q1= C Δt = kj/ o C x ( ) o C = 39.8 kj heat is released by reaction thus q= -q = kj q for run 2 at constant pressure q2= C Δt = kj/ o C x ( ) o C = 34.2 kj heat is released by reaction thus q= -q = kj Missing 5.6 kj???? Expansion work (pressure volume work) w= -PΔV 75 ΔE = q + w Work and heat are simply alternative ways to transfer energy. First law of thermodynamics = law of conservation of energy 76 Some important ideas about energy are summarized in the First Law of Thermodynamics. Energy is neither created nor destroyed in ordinary chemical reactions and physical changes. FIRST LAW OF THERMODYNAMICS The substances involved in the chemical and physical changes that we are studying are called the system. Everything in the system s environment constitutes its surroundings. The universe is the system plus its surroundings. THE UNIVERSE, SYSTEM, SURROUNDINGS

20 The system may be thought of as the part of the universe under investigation. The First Law of Thermodynamics tells us that energy is neither created nor destroyed. Energy is only transferred between the system and its surroundings. Q. The first law of thermodynamics states that energy is a. increased during any process b. decreased during any process c. conserved during any process FIRST LAW OF THERMODYNAMICS FUNDAMMENTAL ASPECT OF THERMODYNAMICS An electron in an atom has potential energy because of the electrostatic force on it that is due to the positively charged nucleus and the other electrons in that atom and surrounding atoms. Potential energy of atom

21 The atomic or molecular level, we can think of each of these as either kinetic or potential energy. The chemical energy in a fuel or food comes from potential energy stored in atoms due to their arrangements in the molecules. ATOMIC LEVEL of energy Many forms of energy can be interconverted and that in chemical processes, chemical energy is converted to heat energy or vice versa. The amount of a heat a process uses (endothermic) or gives off (exothermic) can tell us a great deal about that process. For this reason it is important for us to be able to measure the intensity of the heat. INTENSITY OF HEAT In such reactions, the total energy of the products is lower (for exothermic) or greater (endothermic) than that of the reactants by the amount of energy as a heat released or absorbed. Some initial activation (e.g., by heat)is needed to get these reactions started. This amount of energy is called activation energy ACTIVATION ENERGY 86 21

22 N 2(g) + 3 H 2(g) 2NH 3(g) ΔH 0 = kj The amount of heat shown in thermochemical equation always refers to the reaction for the number of moles of reactants and products specified by the coefficients. 2N 2(g) + 6 H 2(g) 4NH 3(g) ΔH 0 = 2 x( kj ) = kj 1/2N 2(g) H 2 g NH 3 g ΔH 0 = (0.5) kj = kj/ mole THERMOTERMIC REACTION Thermochemical equation H 2(g) + O 2(g) 2H 2 O (l) ΔH 0 = kj H 2 g + ½ O 2(g) H 2 O (l) Reactions that release energy in the form of heat are called exothermic reactions. ΔH 0 =? kj/mole EXOTHERMIC REACTIONS

23 CH 4(g) + 2O 2(g) CO 2(g) + 2H 2 O (l) kj/mol reagents products exothermic reaction at constant pressure. Reactions that absorb energy in the form of heat are called endothermic reactions. Exothermic reaction 91 Endothermic reaction 92 NH 4 NO 3(s) + 26 kj Reagents H 2 O NH 4 NO 3(aq) product ENDOTHERMIC REACTION Energy absorbed or released Relative Energy of reactants & products Sign of H Exothermic Reaction Energy is released. It is a product of the reaction. Reaction vessel becomes warmer. Temperature inside reaction vessel increases. Energy of the reactants is greater than the energy of the products H (reactants) > H (products) H = H (products) - H (reactants) = negative (-ve) Endothermic Reaction Energy is absorbed. It is a reactant of the reaction. Reaction vessel becomes cooler. Temperature inside reaction vessel decreases. Energy of the reactants is less than the energy of the products H (reactants) < H (products) H = H (products) - H (reactants) = positive (+ve)

24 Energy profiles Exothermic Endothermic N H 2 2NH 3 Reversible reaction Energy of the reactants (N 2 & H 2 ) is greater than the energy of the products (NH 3 ). Energy is released. Energy of the reactants (NH 3 ) is less than the energy of the products (N 2 and H 2 ). Energy is absorbed Exothermic processes Endothermic processes making ice cubes formation of snow in clouds condensation of rain from water vapour mixing sodium sulfite and bleach rusting iron burning sugar mixing water and strong acids mixing water with an anhydrous salt crystallizing liquid salts (as in sodium acetate in chemical handwarmers) melting ice cubes conversion of frost to water vapour evaporation of water baking bread cooking an egg producing sugar by photosynthesis mixing water and ammonium nitrate making an anhydrous salt from a hydrate melting solid salts Q. Chemical reactions that absorb heat energy are called. a. exothermic b. eltothermic c. endothermic Q. Electrolysis requires energy to make it work. This means it is... a) an endothermic reaction b) an exothermic reaction c) an eltothermic reaction d) a chemical reaction Author: Fred Senese senese@antoine.frostburg.edu;

25 Q. Which of the following is an endothermic reaction? a) Burning propane in a gas grill b) Photosynthesis c) Baking bread d) Cooking an egg e) Electrolysis of water Q. A. What is the change energy of the sausage after heating, if original energy is 4 kj and 20 kj is added to it? B. What is the total energy content of sauage after heating? a) 16 kj; b) 4 kj; c) 20 kj; d) 24 kj Most chemical reactions and physical changes occur at constant (usually atmospheric) pressure. The quantity of heat transferred into or out of a system as it undergoes a chemical or physical change at constant pressure, q p, is defined as the enthalpy change, H, of the process. ENTHALPY The enthalpy change is equal to the enthalpy or heat content, H, of the substances produced minus the enthalpy of the substances consumed. ENTHALPY - HEAT

26 E H = p V Pressure volume work It is impossible to know the absolute enthalpy (heat content) of a system. The only time when ΔH and ΔE differs by a significant amount is when gases are formed or consumed in a reaction. In such situation we applied ideal gas law and obtain formula for ΔH as follows: Enthalpy is a state function, however, and it is the change in enthalpy in which we are interested. This can be measured for many processes. H = E + n gas RT Difference between energy and enthalpy change EHTHALPY AS STATE FUNCTION Calculate the enthalpy of the ammonium nitrate decomposition. The reaction is and the enthalpies of the three compounds are given in Table 1. ΔH = [ 82+ 2(-242)] [-366] = -36 kj N 2 O H 2 O NH 4 NO 3 products reactant CHANGE OF ENTHALPY

27 If you reverse the previous reaction, the sign of the enthalpy of the reaction is reversed: Δ H = +36 kj nthalpy Endothermic Potential energy products ΔH > 0 reactants Reaction progress Potential energy Exothermic reactants ΔH < 0 products Reaction progress There are two ways of looking at what happens to the enthalpy: If the reaction is exothermic the products have minimum enthalpy and the formation of products (move toward the right) is favourable If the reaction is endothermic the reactants have minimum enthalpy and the formation of products (move toward the right) is unfavourable. In this case the formation of reactants (move toward the left) is favourable

28 Calculate the enthalpy change for the following reaction and classify it as exothermic or endothermic. Standard enthalpies of formation Compound ΔH 0 MgCl 2 (S) -642 kj/mol H 2 O (l) -286 kj/mol MgO (S) HCl (g) -602 kj/mol -92 kj/mol ΔH = [ΔH 0 (MgO) + 2 ΔH 0 (HCl)] [ ΔH 0 (MgCl 2 ) + ΔH 0 (H 2 O)] End of part