Presentation on theme: "Chapter 6 Thermochemistry: The Fire Within This canopy walkway and net in the Peruvian rainforest allow biomass research to be conducted high above the."— Presentation transcript:
Chapter 6 Thermochemistry: The Fire Within This canopy walkway and net in the Peruvian rainforest allow biomass research to be conducted high above the ground. The source of most of our energy is the Sun. Through the process of photosynthesis, solar radiation causes chemical reactions in green plants that store the energy for future use. We make use of the energy captured by plants when we burn fuels. Fossil fuels contain energy that has been stored for thousands of years, but research into alternative fuels is finding ways to make efficient use of plants for fuel. This chapter presents the basic concepts used in research into the energy changes that accompany all chemical reactions. Study of the heat released or required by chemical reactions.
Assignment for Chapter 6 14, 23, 35, 47, 55, 61
Energy, Heat, Enthalpy CH 4 (g)+2O 2 (g) CO 2 (g)+2H 2 O(l)+energy C 6 H 12 O 6 (aq)+O 2 (g) 6CO 2 (g)+6H 2 O(l)+energy Energetics: How energy is transformed and used Bioenergetics: Study of the use of energy in organisms Conservation and transformation of energy Heat Enthalpy
Figure 6.1 The potential energy of a mass, m, is proportional to its height, h, above a surface that is taken to correspond to zero potential energy.
Figure 6.2 Kinetic energy (represented by the height of the dark green bar) and potential energy (the light green bar) can be converted into one another. However, their sum (the total height of the bar) is a constant in the absence of external influences, such as air resistance. A ball thrown up from the ground loses kinetic energy as it slows but gains potential energy. The reverse happens as it falls back to Earth. Conservation of energy: E k +E p =const. for an isolated body Internal Energy U=sum of the kinetic and potential energies of all the atoms and molecules in a sample.
Figure 6.3 In thermodynamics, the world is divided into a system (the object of interest) and the surroundings (everything else). In practice, the surroundings may be a constant-temperature water bath. The arrows represent the energy being transferred between the system and its surroundings. Heat(energy) exchange(transfer) Matter exchange(transfer)
Figure 6.4 We classify systems into one of three kinds, according to their interactions with their surroundings. An open system can exchange matter and energy with its surroundings. A closed system can exchange energy but not matter. An isolated system can exchange neither matter nor energy.
Figure 6.5 The booster rockets on the space shuttle form an open system. The stream of gases produced by the chemical reaction pours out of the engines and moves the rocket. (See Applying Chemistry: Case Study 20.)
Figure 6.6 When we heat a system, we make use of a difference in temperature between it and the surroundings to induce energy to flow through the walls of the system. Heat flows from high temperature to low. Heat is a transfer of energy that occurs as a result of a temperature difference. 1 cal = J
Figure 6.7 When energy leaves a system as a result of a temperature difference between the system and the surroundings, we say that the system has lost energy as heat. This transfer of energy stimulates the thermal motion of molecules in the surroundings. Thermal motion = random molecular motion (in gases, liquids & solids)
Figure 6.8 A system does work when it expands against an external pressure. Here we see a gas that pushes a piston out against a pressure, P. We shall see shortly that the work done is proportional to both the pressure and the change in volume that the system undergoes. Work is a transfer of energy that takes place when an object is moved against an opposing force.
Figure 6.9 When a system expands, it performs work on its surroundings by forcing all the molecules in another object in the surroundings to move in the same direction. Here we see the expansion of a gas raise a weight. The expansion of gases in the cylinders of automobiles do work by pushing on the piston, which turns the gears that move the vehicle.
Three Ways for Changing the Internal Energy of a System Adding matter Supplying energy as heat (thermal motion) Doing work (uniform motion)
Figure 6.10 When we wind a spring, the potential energy of the atoms changes because they are squashed together and repel one another. The internal energy of the spring rises as a result of this increase in potential energy. Another type of potential energy: Interactions between atoms in a system.
Figure 6.11 The internal energy of a system can be changed either by doing work or by heating. The diagram shows that the change in internal energy is positive (U increases) when energy is supplied in either way. When energy leaves the system as heat or work, the internal energy falls and U is negative.
Figure 6.12 The inventor of this elaborate device, the Keely motor, claimed that it could generate high pressures and energies from a small amount of water. However, like other perpetual motion machines, the Keely motor was found to be a fraud. The work was accomplished by compressed air from a hidden source.
Figure 6.13 A reaction does work when it generates a gas. For example, carbon dioxide is formed from the thermal decomposition of calcium carbonate. As the gas is formed, it drives back the surrounding atmosphere. W=?
Figure 6.14 The gas in the piston expands a distance, l, against a pressure, P. The volume increase is l A, where A is the cross-sectional area of the piston.
An Example for Calculating Work A gas expands by 12.0 L against a pressure of 2.0 atm. How much work is done?
Figure 6.15 When a reaction (such as the thermal decomposition of calcium carbonate) takes place in a closed, constant-volume container, the gas fills the container but cannot expand against the surrounding atmosphere. As a result, it does no work on the surroundings.
Figure 6.16 When a system is free to expand against an external pressure, some of the energy supplied to it as heat escapes back into the surroundings as work. As a result, the change in internal energy is less than the energy supplied.
Enthalpy Change in enthalpy of a system is equal to the heat transferred to it at constant pressure
Figure 6.17 The thermite reaction is another highly exothermic reaction—one that can melt the metal it produces. In this reaction, aluminum metal is reacting with iron(III) oxide, Fe 2 O 3, causing a shower of molten iron sparks. In an exothermic reaction, energy is lost as heat, the amount lost depending on the amount of reactants available.
Figure 6.18 The reaction between ammonium thiocyanate, NH 4 SCN, and barium hydroxide octahydrate, Ba(OH) 2 8H 2 O, absorbs a lot of heat and can cause water vapor in the air to freeze on the outside of the beaker. In an endothermic reaction, energy is absorbed as heat. 2NH 4 SCN+ Ba(OH) 2 8H 2 O Ba(SCN) 2 +2NH 3 (g)+10H 2 O (g)
Classroom Exercise In an endothermic reaction, 5 kcal of heat is absorbed and 1.2 L of gas is generated. How much is the change of the internal energy?
Figure 6.19 Heat capacity is an extensive property, so a large object (bottom) has a larger heat capacity than a small object (top) made of the same material.
Specific Heat Capacity
Using Specific Heat Capacity 5.10 g sample of an alloy with specific heat capacity of J/C/g is heated from 24.2 C to C. How much energy is supplied?
Figure 6.20 A pyrotechnics expert attaches fuses containing potassium chlorate to fireworks set up in mortars. The highly exothermic reactions of fireworks do not begin until they are intiated by a fuse. The fuse ignites a mixture of carbon and other reducing agents that react with an oxidizer such as potassium perchlorate.
Figure 6.21 The quantity of heat released or absorbed by a reaction can be measured in this primitive version of a calorimeter. The outer polystyrene cup acts as an extra layer of insulation to ensure that no heat enters or leaves the inner cup. The quantity of heat released or absorbed is proportional to the change in temperature of the calorimeter.
Figure 6.22 A bomb calorimeter. The combustion is initiated with an electrical fuse. Once the reaction has begun, energy is released as heat that spreads through the walls of the bomb into the water. The heat released is proportional to the temperature change of the entire calorimeter assembly.
Calorimetry 50.0 g water at 20 C is mixed with 21 g of iron at 90.2 C. The equilibrium temperature is 23.2 C. Find the specific heat capacity of iron.
Enthalpy H = U + PV Benoit Paul É mile Clapeyron Rudolf Clausius J. Willard Gibbs Heike Kamerlingh Onnes
Thermodynamics related to eating/drinking, so much, more than entire physics …….
H=U+PV How does power come from?
Ice cold water hot water vapor
Vaporization: enthalpy of vaporization
Figure 6.23 Melting (fusion) is an endothermic process. As molecules acquire energy, they begin to struggle past their neighbors. Finally the sample changes from a solid with ordered molecules (left) to a liquid with disordered, mobile molecules (right). Melting/freezing: enthalpy of fusion/freezing
Figure 6.24 The enthalpy change for the reverse of a process is the negative of the enthalpy change for the forward process at the same temperature.
Figure 6.25 The polar ice caps on Mars extend and recede with the seasons. They are solid carbon dioxide and form by direct conversion of the gas to a solid. They disappear by sublimation. Although some water ice is also present in the polar caps, the temperature on Mars never becomes high enough to melt it. On Mars, ice is just another rock.
Figure 6.26 Because enthalpy is a state property, the enthalpy of sublimation at a given temperature can be expressed as the sum of the enthalpies of fusion and vaporization measured at the same temperature. Sublimation: enthalpy of sublimation
Thermochemistry of Some Physical Changes Vaporization: enthalpy of vaporization Melting/freezing: enthalpy of fusion/freezing Sublimation: enthalpy of sublimation
Figure 6.27 The heating curve of water. The temperature of the solid rises as heat is supplied. At the melting point, the temperature remains constant and the heat is used to melt the sample. When enough heat has been supplied to melt all the solid, the temperature of the liquid begins to rise again. A similar pause in the temperature rise occurs at the boiling point.
The Enthalpy of Chemical Change: Reaction Enthalpies Reactants Porducts+heat (exothermic) Reactants+heat Products (endothermic) CH 4 (g)+2O 2 (g) CO 2 (g)+2H 2 O(l)+heat CH 4 (g)+2O 2 (g) CO 2 (g)+2H 2 O(l) ΔH r = kJ/mol (Thermochemical Equation) Molar Reaction Enthalpy (CH 4 )
Figure 6.28 A biophysicist monitors an experimental fermentation chamber in which fuel ethanol is being produced from waste biomass by a genetically engineered strain of bacteria.
+ Many such chemical reactions take place in your body:
Figure 6.29 This diagram shows how the value of the reaction enthalpy depends on the physical states of a product. When water is produced as a vapor rather than as a liquid in the combustion of methane, 88 kJ remains stored in the system for every 2 mol H 2 O produced. CH 4 (g)+2O 2 (g) CO 2 (g)+2H 2 O(g) ΔH r = kJ/mol CH 4 (g)+2O 2 (g) CO 2 (g)+2H 2 O(l) ΔH r = kJ/mol
Figure 6.30 The standard reaction enthalpy is the difference in enthalpy between the pure products, each at 1 atm, and the pure reactants at the same pressure and the specified temperature (which is commonly but not necessarily 25°C). The scheme here is for the combustion of methane. Standard reaction enthalpy: CH 4 (g)+2O 2 (g) CO 2 (g)+2H 2 O(l) ΔH o = kJ/mol
Investigating Matter 6.1 (a) Because fossil fuel reserves are limited, they must be extracted wherever they are found. This platform is used to pump petroleum from beneath the ocean; however, the natural gas accompanying it cannot be easily transported and so is burned off.
Investigating Matter 6.1 (b) An agricultural researcher assesses the growth rate of a seedling. Plant photosynthesis is only about 3% efficient, and conditions that increase this efficiency are actively being sought.
P 4 (s)+6Cl 2 (g) 4PCl 3 (l) 4PCl 3 (l) P 4 (s)+6Cl 2 (g) C 6 H 12 O 6 (aq)+6O 2 (g) 6CO 2 (g)+6H 2 O(l) 6CO 2 (g)+6H 2 O(l) C 6 H 12 O 6 (aq)+6O 2 (g) Reaction Enthalpy of A Reverse Reaction
Figure 6.31 If the overall reaction can be broken down into a series of steps, then the corresponding overall reaction enthalpy is the sum of the reaction enthalpies of the steps on the alternative path. None of the steps need be a reaction that can actually be carried out in the laboratory. Hess’s Law Germain Henri Hess ( )
Using Hess’s Law C(s)+O 2 (g) CO 2 (g) C(s)+(1/2)O 2 (g) CO(g) CO(g)+(1/2)O 2 (g) CO 2 (g) + C(s)+O 2 (g) CO 2 (g)
Using Hess’s Law 3C(s)+4H 2 (g) C 3 H 8 (g) (X) C 3 H 8 (g)+5O 2 (g) 3CO 2 (g)+4H 2 O(l) (A) C(s)+O 2 (g) CO 2 (g) (B) H 2 (g)+(1/2)O 2 (g) H 2 O(g) (C) X=3*B-A+4*C
Figure 6.32 The amount of heat produced or absorbed in a chemical reaction can be determined from the reaction stoichiometry.
The Enthalpy of Chemical Change: Reaction Enthalpies Enthalpy of combustion Enthalpy of formation
CH 4 (g)+2O 2 (g) CO 2 (g)+2H 2 O(g) ΔH c = kJ/mol How much heat is produced by burning g of methane? Enthalpy of Combustion
Figure 6.33 During World War II, fuel was in short supply and all manner of ingenious solutions were sought. However, as we can see from this photograph of a vehicle powered by coal gas (a mixture of carbon monoxide and hydrogen) in London, the low enthalpy density of gases creates storage problems. A modern approach to using gases to power a vehicle can be seen in Applying Chemistry: Case Study 18.
Different Units for Enthalpy Enthalpy Density Specific Enthalpy
Figure 6.34 The range and speed of this electric-powered car depend on the type of battery it uses. For example, metal-hydride devices have a longer range than lead-acid storage batteries. As this driver would agree, recharging is generally a slow process. The cable attached to this car may look like a gasoline hose, but it is actually delivering electricity while its owner waits. The use of hydrogen as fuel could reduce the number of refueling stops.
Standard Enthalpies of Formation The standard enthalpy of formation of an element in its most stable form is 0.
Figure 6.35 The reaction enthalpy can be constructed from enthalpies of formation by imagining the formation of both the reactants and the products from their respective elements. The reaction enthalpy is the difference between the two.
Standard Enthalpies of Formation 4C(s)+6H 2 (g) +O 2 (g) 2C 2 H 5 OH(l) ΔHc= kJ
Classroom Exercise 2C(s)+O 2 (g) 2CO (g) ΔHc= kJ Standard enthalpies of formation of CO?
The standard enthalpy of formation of an element in its most stable form is 0.
Using Standard Enthalpy of Formation
Fun Chemistry: Calorie of Food It is equal to specific combustion enthalpy: kcal/g It is in the sense of average and approximation, depending on the location and growth conditions of the original produces, the genetics of the people who consumes the food … Compound A + O 2 Compound B + Compound C + … ΔHc = calorie It is assumed that the calorie of a compound (protein, carbonhydrate etc) metabolized in body is the same as that when the compound is burnt outside.
Measuring Calories The caloric value of food is the energy produced by combustion of its proteins, carbohydrates and fats. The amount of energy liberated by the catabolism of food in the body is almost the same as the amount liberated when food is burnt outside the body. The energy liberated by catabolic processes is used for maintaining body functions namely digestion, thermoregulation, muscular contraction and nerve impulses conduction. The amount of energy liberated / unit time is the metabolic rate. When food is burnt outside the body, all the energy is liberated as heat. The standard unit of heat energy is the calorie (cal), which is defined as the amount of heat energy necessary to raise the temperature of 1 ml of water by one degree, from 15 to 16 celsius at rest. A slightly different calorie is used in engineering, the international calorie, which equals 1/860 international watt-hour ( J). A large calorie, or kilocalorie, usually referred to simply as a calorie and sometimes as a kilogram calorie, equals 1,000 calories and is the unit used to express the energy-producing value of food in the calculation of diets. The energy released by combustion of foodstuffs outside the body can be measured directly and indirectly. In direct calorimetry method the liberated energy can be measured using a bomb calorimeter. It is a metal vessel surrounded by a water insulated container. The food is ignited by an electric spark. The change in the temperature of water is a measure of the calories produced. In indirect method, the energy production can also be calculated by measuring the amount of oxygen consumed for combustion of food. The amount of oxygen consumed / unit of time is proportionate to the energy liberated. This method of energy estimation is called wet combustion. The caloric value of carbohydrate is 4.1 Kcal /g, protein is 5.65 Kcal/ g and fat is 9.4 Kcal/g.
Case Study 6 (a) Regular exercise not only is good for the metabolism, it can be fun, too, when we make it a part of daily life. When you exercise, you burn the nutrients in your body by speeding up metabolism and spend more calories.
Case Study 6 (b) Energy consumed (in kilojoules per hour) in typical activities: blue for a 70-kg male and pink for a 58-kg female.
Biological systems are wonderful chemicals and wonderful chemical plants
Now, a six-table banquet is served. When enjoying the delicacies, refresh yourself of the meaning of “ calorie ” Dr. Ding told you. Thermodynamics related to eating/drinking, so much, more than entire physics …….