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16.1: Introduction to Energy The source of all energy on earth begins with the sun. Energy is the ability to do work and work is the product of force x distance. There are different forms of energy: radiant energy from the sun, kinetic energy of motion, potential energy of position, chemical stored energy in compounds and nuclear energy stored in the nucleus of the atom. The metric unit for energy is the joule, J (kg•m2/s2) and the Imperial energy unit is the calorie. The conversion factor is 1 cal = 4.184 J. Energy should not be confused with power. Power is the energy transferred per time. The metric unit for power is watt, W (J/s). Sources of Energy and Renewable Energy The numerous traditional sources of energy available for us include oil, natural gas, hydroelectricity, nuclear energy, and coal. The sustainability of the environment has directed society towards seeking alternate sources of energy. The renewable sources of energy include wind, solar, geothermal, ocean, hydro, biomass and hydrogen. Technologies of Renewable Energy The harnessing and the delivery of renewable energy requires the knowledge and tools that enable humans to provide for the demand for energy and yet build an infrastructure that encourages the sustainability of the environment. Wind Energy Wind energy is a renewable energy source that converts the kinetic energy of moving air into the mechanical energy of the generator, which in turn produces electrical energy. The locations of wind turbines range from single units to a series of wind turbines that serve the electrical needs of a wider electrical grid. The dependence and reliability of a good windy area is a very important consideration when deciding the location of the wind turbines. A miscalculation will defeat the purpose of providing a renewable energy source that minimizes the use of the traditional energy sources. The electrical energy produced by wind turbines would be available to other energy conversions. The wind turbines are a very attractive renewable energy source.

Practice Question: Arrange the following types of energy into the proper sequence of energy conversions by the wind turbine: mechanical energy (generator), electrical energy, wind energy (kinetic).

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Solar Energy Harnessing the sun’s radiant energy as a renewable energy source is divided into two sectors: solar thermal and solar photovoltaic. The former relies on the technology that uses solar energy to heat up a medium that will be transported for heating purposes and possibly be turned into electricity, depending on the technology of the design. The later relies on the direct conversion of solar energy into electrical energy. Did You Know? 70% of the energy used in the residential and commercial/institutional buildings sector is used for heating. Canada also has a very large potential for solar energy use and it has excellent solar resources. Since 2007, there are an estimated 544,000 m2 of solar collectors operating in Canada. They are primarily unglazed plastic collectors for pool heating (71%) and unglazed perforated solar air collectors for commercial building air heating (26%), delivering about 627,000 GJ of energy and displacing 38,000 tonnes of CO2 annually. The energy conversions from solar energy depend on the type of solar collector. Photovoltaic cells use solar cells to capture the sun’s energy and make a direct conversion to electrical energy. You might have some photovoltaic cells in your backpack. Some calculators are powered by photovoltaic cells. Solar arrays concentrate the sun’s rays onto a focal point where a pipe is carrying water. The shape of the array is parabolic. The concentrated energy is sufficient enough to produce significant change in the thermal energy of the water. Hydroelectric Power The harnessing of the gravitational potential energy found in water in terms of large and small scales, is another avenue for producing electricity as opposed to fossil-burning power plants. The traditional large-scale hydroelectric power plants, like Niagara Falls and the James Bay units, are the more commonly known sources of hydroelectric power. The large units serve a very wide and expansive energy grid. There is a large dependency on these units by society. Smaller units are now available to areas where there is a need for electrical power and the availability for deploying such units. Practice Question: What by-product of fossil-burning power plants would you find contributing to the greenhouse emissions?

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Hydropower Small hydropower is a clean, renewable, and predictable energy source. Small hydropower generation systems produce electricity by converting mechanical energy in the running water into electric energy in a way similar to larger traditional hydroelectric systems. In Canada, small hydro generally refers to hydroelectric projects with between 1 and 50 megawatts (MW) in installed capacity. Small hydro systems can either be connected to the grid and provide power to the grid or they can be used for independent and stand-alone applications in isolated remote areas. Small hydro is one of the best alternatives to the highly polluting and very costly diesel generation that currently provides electric energy in most remote communities across Canada. The potential energy that could be produced by small hydropower technologies is estimated to be 15,000 MW, and the current installed small hydro capacity is approximately 3,400 MW. By using both the existing state-of-the-art and the emerging technologies, the vast small hydro potential could be developed to help Canada meet its future energy needs and bring about more environmental and socio-economical benefits. The following diagram shows a cross-section of a typical hydroelectric power generating station, followed by a brief summary of the energy conversions resulting from the falling water.

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Hydroelectric power generation The very large potential energy of the water due to its height and mass, converts to the kinetic energy of the moving of the water. The generators convert the kinetic energy into the mechanical energy of the turbines. Electrical energy is the final useable energy product that is available to the electrical grid serving the needs of consumers and businesses. The large scale developments of these units depend on the economics of the site and the return in useable and efficient energy. The effects on the environment are pronounced in today’s world and this is another serious consideration in their construction. Hydrogen Power The potential for using hydrogen as a power supply is far reaching and probably to many peoples’ surprise, the uses of hydrogen power extends from the fuel cells in cars to a portable cell phone Hydrogen makes up 3/4 of the known universe. However, it almost never appears alone and must be separated from its chemical compounds to be useful. This can be done through the reforming of hydrocarbons (oil and natural gas), or using electricity to split the hydrogen out of water (electrolysis). Hydrogen is an energy carrier - like gasoline or electricity - not an energy source like oil or coal. Unlike electricity, however, hydrogen can be stored. As a fuel, hydrogen's only byproduct is water, making it a very clean and environmentally-friendly alternative fuel. Hydrogen can be used as a transportation fuel or in stationary power applications, using fuel cells of varying sizes and types. A fuel cell is a device that converts chemical energy into electrical energy. Hydrogen combines with oxygen to produce electricity. The only emissions are water and heat. The three fuel cell market segments are: portable applications such as cell phones and laptops; stationary applications for residential and commercial energy sources such as heating, cooling and electricity, and mobile applications such as automobiles and forklifts Biomass The harnessing of energy from biomass is another sector of the renewable energy initiatives in the province of Ontario. The sources of biomass are shown in the following diagram: On the subject of biomass energy, the Ontario Ministry of Energy and Infrastructures states the following: Biomass from renewable plant and animal materials can be used to produce heat or power. Burning biomass to produce power results in substantially fewer harmful emissions when compared to traditional sources of power generation. Within Ontario there are extensive opportunities to make use of biomass materials from landfill sites, agricultural and livestock operations, and the forest industry. Ontario Regulation 232/98 under the Environmental Protection Act requires the collection of landfill gas for new or expanding landfill sites larger than three million cubic metres or 2.5 million tonnes. The production of energy from methane

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derived from animal manure wastes can serve to mitigate other disposal and surface or groundwater contamination concerns. Promising technological developments related to these fuel/energy technologies, such as anaerobic digestion technology could also have positive environmental and economic spin-offs for Ontario. The search for energy, its transmittance, storage and efficient use alongside the goal of ensuring a sustainable environment are of constant concern in today’s society. The efforts to minimize the effects on the environment and maximize the efficiencies of coal-based electric power plants and the catalytic converter for the automobile are the two examples of this complex relationship between sustainability and the demand for energy. Coal-Fired Power Plants Traditional coal-fired power plants are based on “yesterday’s” technology. The efficient conversion of energy to useful energy minimizes the effects on the environment. However, technological advancements have been incorporated to reduce emissions that contribute to greenhouse emissions.

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Inquiry Questions

1. Show the energy conversions, in order, in a coal-fired power plant.

2. What criteria for efficiency was the old technology based on?

3. What new conditions are proposed to increase energy efficiency and minimize carbon dioxide emissions?

Catalytic Converter The combustion of nitrogen gas yields nitrogen monoxide. The incomplete combustion of hydrocarbons can also yield carbon monoxide. These two reactions occur in the engine of a car. Nitrogen monoxide will oxidize in the atmosphere to produce nitrogen dioxide, and this compound is a contributing factor in the brownish haze found in air pollution. The catalytic converter converts the unwanted products in the exhaust into nitrogen gas, water and carbon dioxide. This process, too, is not 100% efficient. However, there is constant research and development to improve the design of the catalytic converter. Inquiry Questions:

1. The interior of the converter appears to have many sites on its surface where a reaction could take place. What physical property is considered in the design of a catalytic converter?

2. What condition was a major factor in the proper operation of a catalytic converter?

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16.2: Changes in Matter Thermochemistry is the study of the energy changes that accompany these transformations. Changes that occur in matter may be classified as physical, chemical, or nuclear, depending on whether a change has occurred in the arrangements of the molecules, their electronic structure, or the nuclei of the atoms involved. Whether ice melts, iron rusts, or an isotope used in medical therapy undergoes radioactive decay, changes occur in the energy of chemical substances. Hydrogen may undergo a physical, chemical, or nuclear change: a.) Physical: Hydrogen boils at 252°C (or only about 20°C above absolute zero): H2(l) → H2(g) b.) Chemical: Hydrogen is burned as fuel in the space shuttle’s main engines: 2H2(g) + O2(g) → 2H2O(l) c.) Nuclear: Hydrogen undergoes nuclear fusion in the Sun, producing helium: H + H → He Heat and Energy Changes Both physical and chemical changes are involved in the operation of an oxyacetylene torch to weld metals together. A chemical reaction, which involves ethyne (or acetylene) and oxygen as reactants, produces carbon dioxide gas, water vapour, and considerable energy .This energy is released to the surroundings as thermal energy, a form of kinetic energy that results from the motion of molecules. The result is a physical change — the melting of the metal — when the increased vibration of metal particles causes them to break out of their ordered solid pattern. When you are studying such transfers of energy, it is important to distinguish between the substances undergoing a change, called the chemical system, and the system’s environment, called the surroundings. A system is often represented by a chemical equation. For the burning of ethyne, the equation is:

2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(g) + energy The surroundings in this reaction would include anything that could absorb the thermal energy that has been released, such as metal parts, the air, and the welder’s protective clothing. When the reaction occurs, heat, q, is transferred between substances. (An object possesses thermal energy but cannot possess heat.) When heat transfers between a system and its surroundings, measurements of the temperature of the surroundings are used to classify the change as exothermic or endothermic. In exothermic changes, energy is released from the system, usually causing an increase in the temperature of the surroundings. In endothermic changes, energy is absorbed by the system, usually causing a decrease in the temperature of the surroundings. The acetylene torch reaction is clearly an exothermic reaction because heat flows into the surroundings. Chemical potential energy in the system is converted to heat energy, which is transferred to the surroundings and used to increase

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the thermal energy of the molecules of metal and air. Since the molecules in the surroundings have greater kinetic energy, the temperature of the surroundings increases measurably. A chemical reaction that produces a gas in a solution in a beaker is described as an open system, since both energy and matter can flow into or out of the system. An isolated system is an ideal system in which neither matter nor energy can move in or out closed system one in which energy can move in or out, but not matter. Measuring Energy Changes: Calorimetry When methane reacts with oxygen in a lab burner, enough heat is transferred to the surroundings to increase the temperature and even to cause a change of state. How is this amount of heat measured? The experimental technique is called calorimetry and it depends on careful measurements of masses and temperature changes. When a fuel like methane burns, heat is transferred from the chemical system into the surroundings (which include the water in the beaker). Different substances vary in their ability to absorb amounts of heat. These three factors — mass (m),temperature change (∆T),and type of substance — are combined in an equation to represent the quantity of heat (q) transferred: q=mc∆T where c is the specific heat capacity, the quantity of heat required to raise the temperature of a unit mass (e.g., one gram) of a substance by one degree Celsius or one kelvin.For example,the specific heat capacity of water is 4.18 J/(g•°C). Inquiry Questions:

1. When 600.0 mL of water in an electric kettle is heated from 20.0°C to 85.0°C to make a cup of tea, how much heat flows into the water?

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2. What would the final temperature be if 250.0 J of heat were transferred into 10.0 g of methanol initially at 20.0°C?

Heat Transfer and Enthalpy Change Chemical systems have many different forms of energy, both kinetic and potential. These include the kinetic energies of:

x moving electrons within atoms; x the vibration of atoms connected by chemical bonds; and x the rotation and translation of molecules that are made up of these atoms.

More importantly, they also include:

x the nuclear potential energy of protons and neutrons in atomic nuclei; and x the electronic potential energy of atoms connected by chemical bonds.

Chemists usually study the enthalpy change, or the energy absorbed from or released to the surroundings when a system changes from reactants to products. An enthalpy change is given the symbol ∆H, and can be determined from the energy changes of the surroundings. A useful assumption that will be applied in more detail later in this chapter is that the enthalpy change of the system equals the quantity of heat that flows from the system to its surroundings, or from the surroundings to the system.

∆𝐻 = ± 𝑞

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This idea is consistent with the law of conservation of energy — energy may be converted from one form to another, or transferred from one set of molecules to another, but the total energy of the system and its surroundings remains the same. For example, consider the reaction that occurs when zinc metal is added to hydrochloric acid in a flask:

Zn(s) + 2HCl(aq) → H2(g) + ZnCl2(aq) Some of the chemical potential energy in the system is converted initially to increased kinetic energy of the products. Eventually, through collisions, this kinetic energy is transferred to particles in the surroundings. The enthalpy change in the system is equal to the heat released to the surroundings. We can observe this transfer of energy, and can measure it by recording the increase in temperature of the surroundings (which include the solvent water molecules, the flask, and the air around the flask). We can observe enthalpy changes during phase changes, chemical reactions, or nuclear reactions. Most of the energy is produced in a nuclear change than in a chemical change, and in a chemical change than in a physical change. Types of Enthalpy Changes Physical Changes

x Energy is used to overcome or allow intermolecular forces to act. x Fundamental particles remain unchanged at the molecular level. x Temperature remains constant during changes of state

(e.g., water vapour sublimes to form frost: H2O(g) → H2O(s) + heat). x Temperature changes during dissolving of pure solutes

(e.g., potassium chloride dissolves: KCl(s) + heat → KCl(aq)). x Typical enthalpy changes are in the range ∆H = 100 – 102 kJ/mol.

Chemical Changes

x Energy changes overcome the electronic structure and chemical bonds within the particles (atoms or ions).

x New substances with new chemical bonding are formed (e.g., combustion of propane in a barbecue: C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g) + heat); (e.g., calcium reacts with water: Ca(s) 2H2O(l) → H2(g) + Ca(OH)2(aq) + heat).

x Typical enthalpy changes are in the range ∆H = 102 – 104 kJ/mol. Nuclear Changes

x Energy changes overcome the forces between protons and neutrons in nuclei. x New atoms, with different numbers of protons or neutrons, are formed

(e.g., nuclear decay of uranium-238: 𝑈  → 𝐻𝑒 +  𝑇ℎ  + ℎ𝑒𝑎𝑡). x Typical enthalpy changes are in the range ∆H = 1010 – 1012 kJ/mol. The magnitude of the energy

change is a consequence of Einstein’s equation (E=mc2, c = speed of light 3.0x102 m/s ).

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16.3: Molar Enthalpies When we write an equation to represent changes in matter, the chemical symbols may represent individual particles but usually they represent numbers of moles of particles. Thus, the thermochemical equation: H2(g)+ ½O2(g) Æ H2O(g) + 241.8 kJ represents the combustion reaction of 1 mol of hydrogen with 0.5 mol of oxygen to form 1 mol of water vapour. The enthalpy change per mole of a substance undergoing a change is called the molar enthalpy and is represented by the symbol ∆Hx, where x is a letter or a combination of letters to indicate the type of change that is occurring. Thus, the molar enthalpy of combustion of hydrogen is ∆Hcomb = – 241.8 kJ/mol Note the negative sign in the value of ∆H. Changes in matter may be either endothermic or exothermic. The following sign convention has been adopted. Enthalpy changes for exothermic reactions are given a negative sign. Enthalpy changes for endothermic reactions are given a positive sign. Some Molar Enthalpies of Reaction (∆Hx)

Type of molar enthalpy

solution (∆Hsol): NaBr(s) → Na+(aq) + Br– (aq) combustion (∆Hcomb): CH4(g) + 2O2(g) → CO2(g) + H2O(l) vaporization (∆Hvap): CH3OH(l) → CH3OH(g) freezing (∆Hfr): H2O(l) → H2O(s) neutralization (∆Hneut)*: 2NaOH(aq) + H2SO4(aq) → 2Na2SO4(aq) + 2H2O(l) neutralization (∆Hneut)*: NaOH(aq) + 1/2H2SO4(aq) → 1/2Na2SO4(aq) + H2O(l) formation (∆Hf)**: C(s) + 2H2(g) + 1/2O2(g) → CH3OH(l)

* Enthalpy of neutralization can be expressed per mole of either base or acid consumed. We can express the molar enthalpy of a physical change, such as the vaporization of water, as follows:

H2O(l) + 40.8 kJ → H2O(g) What we may think of as the change in potential energy in the system, the molar enthalpy of vaporization for water, is

∆Hvap = 40.8 kJ/mol Molar enthalpy values are obtained empirically and are listed in reference books in tables such as the one shown below.

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Molar Enthalpies for Changes in State of Selected Substances Chemical Name Formula Molar Enthalpy of Fusion

(kJ/mol) Molar Enthalpy of Vaporization

(kJ/mol)

sodium Na 2.6 101

chlorine Cl2 6.40 20.4

sodium chloride NaCl 28 171

water H2O 6.03 40.8

ammonia NH3 — 1.37

freon-12 CCl2F2 — 34.99

methanol CH3OH — 39.23

ethylene glycol C2H4(OH)2 — 58.8

The amount of energy involved in a change (the enthalpy change ∆H, expressed in kJ) depends on the quantity of matter undergoing that change. This is logical: twice the mass of ice will require twice the amount of energy to melt. To calculate an enthalpy change ∆H for some amount of substance other than a mole, you need to obtain the molar enthalpy value ∆Hx from a reference source and then use the formula ∆H = n∆Hx. Inquiry Problems

1. A common refrigerant is alternately vaporized in tubes inside a refrigerator, absorbing heat, and condensed in tubes outside the refrigerator, releasing heat. This results in energy being transferred from the inside to the outside of the refrigerator. The molar enthalpy of vaporization for the refrigerant is 34.99 kJ/mol. If 500.0 g of the refrigerant is vaporized, what is the expected enthalpy change ∆H?

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2. What amount of ethylene glycol would vaporize while absorbing 200.0 kJ of heat?

Calorimetry of Physical Changes Studying energy changes requires an isolated system, that is, one in which neither matter nor energy can move in or out. Carefully designed experiments and precise measurements are also needed. Two nested disposable polystyrene cups are a fairly effective calorimeter for making such measurements. When we investigate energy changes we base our analysis on the law of conservation of energy: the total energy change of the chemical system is equal to the total energy change of the surroundings.

∆𝐻 = ± 𝑞 There are three simplifying assumptions often used in calorimetry:

x no heat is transferred between the calorimeter and the outside environment; x any heat absorbed or released by the calorimeter materials, such as the container, is negligible; & x a dilute aqueous solution is assumed to have a density and specific heat capacity equal to that of

pure water (1.00 g/mL and 4.18 J/g•°C or 4.18 kJ/kg • °C).

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Inquiry Questions

1. In a calorimetry experiment, 7.46 g of potassium chloride is dissolved in 100.0 mL (100.0 g) of water at an initial temperature of 24.1°C. The final temperature of the solution is 20.0°C. What is the molar enthalpy of solution of potassium chloride?

2. What mass of lithium chloride must have dissolved if the temperature of 200.0 g of water increased by 6.0°C? The molar enthalpy of solution of lithium chloride is –37 kJ/mol.

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Calorimetry of Chemical Changes Chemical reactions that occur in aqueous solutions can also be studied using a polystyrene calorimeter. The chemical system usually involves aqueous reactant solutions that are considered to be equivalent to water. The assumptions and formulas applied are identical to those used in the analysis of energy changes during state changes and dissolving. When aqueous solutions of acids and bases react, they undergo a neutralization reaction. For example, potassium hydroxide and hydrobromic acid solutions react to form water and aqueous potassium bromide:

KOH(aq) + HBr(aq) → H2O(l) + KBr(aq) The molar enthalpy of reaction for systems such as this is sometimes called the heat of neutralization, or enthalpy of neutralization.

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16.4: Representing Enthalpy Changes it is usually not obvious whether a chemical change will absorb or release energy, so, when we are discussing thermochemical reactions, we must indicate this information clearly. The equations we use to do this are called thermochemical equations. Remember: Endothermic enthalpy changes are reported as positive values; and exothermic enthalpy changes are reported as negative values. When water decomposes, the system gains energy from the surroundings and so the molar enthalpy is reported as a positive quantity to indicate an endothermic change: H2O(g) Æ H2(g)+ ½O2(g) ∆Hdecomp = +285.8 kJ/mol H2O The law of conservation of energy implies that the reverse process (combustion of hydrogen) has an equal and opposite energy change H2(g) + ½O2(g) Æ H2O(g) ∆Hcomb = –285.8 kJ/mol H2 The sign convention represents the change from the perspective of the chemical system itself, not from that of the surroundings. An increase in the temperature of the surroundings implies a decrease in the enthalpy of the chemical system, because the change was exothermic. We can communicate the energy changes, obtained from these empirical studies, in four different ways. Here is an example of all four methods when burning methanol

x Method 1: by including an energy value as a term in the thermochemical equation

CH3OH(l) + O2(g) Æ CO2(g) +2H2O(g) +726 kJ

x Method 2: by writing a chemical equation and stating its enthalpy change

CH3OH(l) + O2(g) Æ CO2(g) + 2H2O(g) ∆H= –726 kJ

x Method 3: by stating the molar enthalpy of a specific reaction ∆Hcombustion or ∆Hc = –726 kJ/mol CH3OH

x Method 4: by drawing a chemical potential energy diagram

Method 1: Thermochemical Equations with Energy Terms If a reaction is endothermic, it requires a certain quantity of energy to be supplied to the reactants. If a reaction is exothermic, it requires a certain quantity of energy to be supplied to the products: H2O(l) + 285.8 kJ Æ H2(g) + ½ O2(g) If a reaction is exothermic, energy is released as the reaction proceeds and is listed along with the products: Mg(s) + ½ O2(g) Æ MgO(s) + 601.6 kJ

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Inquiry Questions: 1. Write a thermochemical equation to represent the exothermic reaction that occurs when two

moles of butane burn in excess oxygen gas. The molar enthalpy of combustion of butane is –2871 kJ/mol.

2. Write a thermochemical equation to represent the dissolving of one mole of silver nitrate in water. The molar enthalpy of solution is + 22.6 kJ/mol.

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Method 2: Thermochemical Equations with ∆H Values  A second way to describe the enthalpy change in a reaction is to write a balanced chemical equation and then the ∆H value beside it, making sure that ∆H is given the correct sign. Note: Units for the enthalpy change are kilojoules (not kJ/mol), because the enthalpy change applies to the reactants and products as written, with the numbers of moles of reactants and products given in the equation. CO(g) + H2(g) → CH3OH(l) ∆H= –128.6 kJ ½CO(g) + H2(g) → ½CH3OH(l) ∆H= –64.3 kJ Inquiry Questions:

1. Sulfur dioxide and oxygen react to form sulfur trioxide. The molar enthalpy for the combustion of sulfur dioxide, ∆Hcomb, in this reaction is –98.9 kJ/mol SO2. What is the enthalpy change for this reaction?

2. Write a thermochemical equation, including a ∆H value, to represent the exothermic reaction between xenon gas and fluorine gas to produce solid xenon tetrafluoride, given that the reaction produces 251 kJ/mol of Xe reacted.

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Method 3: Molar Enthalpies of Reaction As you have seen in the previous section, molar enthalpies are convenient ways of describing the energy changes involved in a variety of physical and chemical changes. In each case, one mole of a particular reactant or product is specified. For example, the enthalpy change involved in the dissolving of one mole of solute is called the molar enthalpy of solution and can be symbolized by ∆Hsol. In the table below, the substance under consideration in each reaction is highlighted in red. A molar enthalpy that is determined when the initial and final conditions of the chemical system are at SATP is called a standard molar enthalpy of reaction. The symbol ∆Hx° distinguishes standard molar enthalpies from molar enthalpies, ∆Hx, which are measured at other conditions of temperature and pressure.Standard molar enthalpies allow chemists to create tables to compare enthalpy values. Some Molar Enthalpies of Reactions

Type of molar enthalpy Example of change

solution (∆Hsol) NaBr(s) → Na+(aq) + Br–(aq)

combustion (∆Hcomb) CH4(g) + 2O2(g) → CO2(g) + H2O(l)

vaporization (∆Hvap) CH3OH(l) → CH3OH(g)

freezing (∆Hfr) H2O(l) → H2O(s)

neutralization (∆Hneut)* 2NaOH(aq) + H2SO4(aq) → 2Na2SO4(aq) + 2H2O(l)

neutralization (∆Hneut)* NaOH(aq) + ½H2SO4(aq) → ½Na2SO4(aq) + 2H2O(l)

formation (∆Hf) C(s) + 2H2(g) + ½O2(g) → CH3OH(l) * Enthalpy of neutralization can be expressed per mole of either base or acid consumed For an exothermic reaction, the standard molar enthalpy is measured by taking into account all the energy required to change the reaction system from SATP, in order to initiate the reaction, and all the energy released following the reaction, as the products are cooled to SATP. For example, the standard molar enthalpy of combustion of methanol is ∆H°c = –726 kJ/mol CH3OH This quantity takes into account the energy input to initiate the reaction, the burning of 1 mol of methanol in oxygen to produce 1 mol CO2(g) and 2 mol H2O(g), then the energy released as the products are cooled to SATP. Methanol is produced industrially by the high-pressure reaction of carbon monoxide and hydrogen gases: CO(g) + 2H2(g) → CH3OH(l) Chemists have determined the standard molar enthalpy of reaction for methanol in this reaction, ∆H°r, to be –128.6 kJ/mol CH3OH. To describe the reaction fully, we would write the thermochemical equation CO(g) + 2H2(g) → CH3OH(l) ∆H°r = –128.6 kJ/mol CH3OH The symbol for the molar enthalpy of reaction uses the subscript “r” to refer to the reaction under consideration, with the stated number of moles of reactants and products. Since two moles of hydrogen

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are consumed as 128.6 kJ of heat are produced, the standard molar enthalpy of reaction in terms of hydrogen could be described as half the above value, or –64.3 kJ/mol H2. Inquiry Questions:

1. Write an equation whose energy change is the molar enthalpy of combustion of propanol (C3H7OH).

2. Write an equation whose enthalpy change is the molar enthalpy of reaction of calcium with hydrochloric acid to produce hydrogen gas and calcium chloride solution.

Method 4: Potential Energy Diagrams Potential energy is stored or released as the positions of the particles change, just as it is when a spring is stretched and then released. As bonds break and re-form and the positions of atoms are altered, changes occur in potential energy. As you have seen before, the potential energy change in the system is equivalent to the heat transferred to or from the surroundings. The vertical axis on the diagram represents the potential energy of the system. Since the reactants are written on the left and the products on the right, the horizontal axis is sometimes called a reaction coordinate or reaction progress. In an exothermic change, the products have less potential energy than the reactants: energy is released to the surroundings as the products form.

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In an endothermic change, the products have more potential energy than the reactants: energy is absorbed from the surroundings. Neither of the axes is numbered; only the numerical change in potential energy (enthalpy change, ∆H) of the system is shown in the diagrams. 

Recap of the 4 Methods:

1. Thermochemical Equations with Energy Terms 2. Thermochemical Equations with ∆H Values 3. Molar Enthalpies of Reaction 4. Potential Energy Diagrams

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16.5: Hess’s Law of Additivity of Reaction Enthalpies The reaction to produce carbon monoxide from its elements is impossible to measure with a calorimeter because the combustion of carbon produces both carbon dioxide and carbon monoxide simultaneously.

In this potential energy diagram, nitrogen gas and oxygen gas combine to form nitrogen dioxide, but there are two different paths to reach the products. In one path, nitrogen (N2) and oxygen (O2) gases react to form nitrogen monoxide (NO), a reaction for which ∆H = +90 kJ. Then, nitrogen monoxide and more oxygen react to form nitrogen dioxide (NO2) gas, a reaction for which ∆H = – 56 kJ. In the other path, nitrogen (N2) and oxygen (O2) gases react directly to form nitrogen dioxide (NO2) gas. In both cases, the overall enthalpy change, ∆H = +34 kJ, is the same. Hess’s Law The value of the ∆H for any reaction that can be written in steps equals the sum of the values of H for each of the individual steps. Another way to state Hess’s law is: If two or more equations with known enthalpy changes can be added together to form a new “target” equation, then their enthalpy changes may be similarly added together to yield the enthalpy change of the target equation.

∆𝐻 = ∑∆𝐻 ∆Htarget = ∆H1 + ∆H2 + ∆H3 …

Hess’s discovery allowed chemists to determine the enthalpy change of a reaction without direct calorimetry, using two familiar rules for chemical equations and enthalpy changes:

x If a chemical equation is reversed, then the sign of ∆H changes. x If the coefficients of a chemical equation are altered by multiplying or dividing by a constant

factor, then the ∆H is altered in the same way

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Sample Problem: What is the enthalpy change for the formation of two moles of nitrogen monoxide from its elements? Answer: N2(g) + O2(g) → 2NO(g) ∆H° = ? This reaction, which may be called the target equation to distinguish it clearly from the other equations, is difficult to study calorimetrically since the combustion of nitrogen produces nitrogen dioxide as well as nitrogen monoxide. However, we can measure the enthalpy of complete combustion in excess oxygen (to nitrogen dioxide) for both nitrogen and nitrogen monoxide by calorimetry. Consider the following two known reference equations:

(1) N2(g) + O2(g) → 2NO(g) ∆H1° = +34 kJ

(2) NO(g) + ½O2(g) → NO2(g) ∆H2° = –56 kJ If we work with these two equations, which may be called known equations, and then add them together, we obtain the chemical equation for the formation of nitrogen monoxide. The first term in the target equation for the formation of nitrogen monoxide is one mole of nitrogen gas. We therefore need to double equation (1) so that N2(g) will appear on the reactant side when we add the equations. However, from equation (2) we want 2 mol of NO(g) to appear as a product, so we must reverse equation (2) and double each of its terms (including the enthalpy change). Effectively, we have multiplied known equation (1) by +2, and multiplied known equation (2) by –2.

2✕ (1): N2(g) + O2(g) → 2NO(g) ∆H1° = 2(+34 kJ)

–2✕ (2): 2NO2(g) → 2NO(g) + O2(g) ∆H2° = –2(–56 kJ) Note that the sign of the enthalpy change in equation (2) will change, since the equation has been reversed. Now add the reactants, products, and enthalpy changes to obtain a net reaction equation. Note that 2NO2(g) can be cancelled because it appears on both sides of the net equation. Similarly, O2(g) can be cancelled from each side of the equation, yielding the target equation: N2(g) + 2O2(g) + 2NO2(g) → 2NO2(g) + 2NO(g) + O2(g) which simplifies to: N2(g) + O2(g) → 2NO(g) Now we can apply Hess’s law: If the known equations can be added together to form the target equation, then their enthalpy changes can be added together. ∆H° = (2 ✕ 34) kJ + (–2 ✕(–56)) kJ = +68 kJ + 112kJ ∆H°  = +180 kJ The enthalpy change for the formation of two moles of nitrogen monoxide from its elements is 180 kJ. When manipulating the known equations, you should check the target equation and plan ahead to ensure that the substances end up on the correct sides and in the correct amounts.

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Inquiry Questions: 1. What is the enthalpy change for the formation of one mole of butane gas from its elements?

The reaction is: 4C(s) + 5H2(g) → C4H10(g) ∆H° = ?

The following known equations, determined by calorimetry, are provided: (1): C4H10(g) + O2(g) → 4CO2(g) + 5H2O(g) ∆H1° –2657.4 kJ (2): C(s) + O2(g) → CO2(g) ∆H2° = –393.5 kJ (3): 2H2(g) + O2(g) → 2H2O(g) ∆H3°= –483.6 kJ

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2. Determine the enthalpy change involved in the formation of two moles of liquid propanol. 6C(s) + 8H2(g) + O2(g) → 2C3H7OH(l)

The standard enthalpies of combustion of propanol, carbon, and hydrogen gas at SATP are –2008, –394, and –286 kJ/mol, respectively.

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Multistep Energy Calculations Several energy calculations might be required, involving a combination of energy change definitions such as:

x Heat flow: q=mc∆T x Enthalpy changes: ∆H = n∆Hr x Hess’s Law:  ∆𝐻 = ∑∆𝐻

In these multi-step problems, ∆H is often found by using standard molar enthalpies or Hess’s law and then equated to the transfer of heat, q. If we know the enthalpy change of a reaction and the quantity of reactant or product, we can predict how much energy will be absorbed or released. Inquiry Questions:

1. In the Solvay process for the production of sodium carbonate (or washing soda), one step is the endothermic decomposition of sodium hydrogen carbonate:

2NaHCO3(s) + 129.2 kJ → Na2CO3(s) + CO2(g) + H2O(g)

What quantity of chemical energy, ∆H, is required to decompose 100.0 kg of sodium hydrogen carbonate?

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2. How much energy can be obtained from the roasting of 50.0 kg of zinc sulfide ore?

ZnS(s) + O2(g) → ZnO(s) + SO2(g)

You are given the following thermochemical equations. (1) ZnO(s) → Zn(s) + ½ O2(g) ∆H1° = 350.5 kJ (2) S(s) + O2(g) → SO2(g) ∆H2° = –296.8 kJ (3) ZnS(s) → Zn(s) + S(s) ∆H3° = 206.0 kJ

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16.6: Standard Enthalpies of Formation Calorimetry and Hess’s law are two ways of determining enthalpies of reaction. A third method uses tabulated enthalpy changes (standard enthalpies of formation) for a special set of reactions called formation reactions, in which compounds are formed from their elements. For example, the formation reaction and standard enthalpy of formation for carbon dioxide are: C(s) + O2(g) → CO2(g) ∆Hf˚ = –393.5 kJ/mol Both the elements on the left side of the equation are in their standard states — their most stable form at SATP (25°C and 100 kPa). Note also that the units of standard enthalpies of formation are kJ/mol because they are always stated for that quantity of substance. Writing Formation Equations Formation equations are always written for one mole of a particular product, which may be in any state or form, but the reactant elements must be in their standard states. For example, the standard states of most metals are monatomic solids (Mg(s), Ca(s), Fe(s), Au(s), Na(s)), some nonmetals are diatomic gases (N2(g), O2(g), H2(g)), and the halogen family shows a variety of states (F2(g), Cl2(g), Br2(l), I2(s)). The periodic table at the back of this text identifies the states of elements. Inquiry Questions:

1. Write the formation equation for liquid ethanol.

2. What is the formation equation for liquid carbonic acid?

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Using Standard Enthalpies of Formation Consider the equation for the formation of hydrogen gas: H2(g) → H2(g) The product and reactant are the same, so there is no change in the enthalpy of the system. This observation can be generalized to all elements in their standard states: ∆Hf˚ for Elements  The standard enthalpy of formation of an element already in its standard state is zero. Thus, the standard enthalpies of formation of, for example, Fe(s), O2(g), and Br2(l) are all zero. Standard molar enthalpies of formation give us a means of comparing the stabilities of substances. For example, the element carbon exists in two solid forms at SATP: diamond, used in jewelry and mining drill bits; and graphite, the black substance used in pencil “leads” and composite plastics. Graphite is the more stable form of carbon at SATP, so ∆H˚f(graphite) = 0 kJ/mol Diamond is slightly less stable at SATP and has a greater potential energy than graphite (as shown in the graph). For the formation of diamond: C(graphite) → C(diamond) ∆H˚f(diamond)= +1.9 kJ/mol You have seen that Hess’s law may be applied to a set of known equations to find an unknown enthalpy change. We can apply this problem-solving method to predict the energy changes for many reactions.

∆H = Σn∆H°f(products) – Σn∆H°f(reactants)

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Inquiry Questions: 1. What is the thermochemical equation for the reaction of lime (calcium oxide) and water?

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2. The main component in natural gas used in home heating or laboratory burners is methane. What is the molar enthalpy of combustion of methane fuel?

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3. The standard enthalpy of combustion of benzene (C6H6(l)) to carbon dioxide and liquid water is –3273 kJ/mol. What is the standard enthalpy of formation of benzene, given the tabulated values for carbon dioxide (∆H˚f = –393.5 kJ) and liquid water (∆H˚f = –285.8 kJ)?

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Multistep Energy Calculations Using Standard Enthalpies of Formation Chemical engineers frequently need to do calculations in which they determine the heats produced by an internal-combustion engine. Such calculations involve bringing together many of the problem-solving skills that you are developing. Two key relationships that you applied were:

I. enthalpy change in the system heat transferred to/from the surroundings ∆H = q and II. ∆H = n∆H˚r

Inquiry Questions:

1. When octane burns in an automobile engine, heat is released to the air and to the metal in the car engine, but a significant portion is absorbed by the liquid in the cooling system—an aqueous solution of ethylene glycol. What mass of octane is completely burned to cause the heating of 20.0 kg of aqueous ethylene glycol automobile coolant from 10.0°C to 70.0°C? The specific heat capacity of the aqueous ethylene glycol is 3.5 J/(g •°C). Assume water is produced as a gas and that all the heat flows into the coolant.

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2. One way to heat water in a home or cottage is to burn propane. If 3.20 g of propane burns, what temperature change will be observed if all of the heat from combustion transfers into 4.0 kg of water?