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Thermochemistry - PowerPoint PPT Presentation

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Thermochemistry. Energy. Energy is necessary for all life. The study of energy and it transformations is known as thermodynamics.

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Energy l.jpg

  • Energy is necessary for all life.

  • The study of energy and it transformations is known as thermodynamics.

  • In chemical reactions we study that aspect of thermodynamics that involves the relationships between chemical reactions and energy changes involving heat. This relationship is called thermochemistry.

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The nature of energy

  • Energy is defined as the capacity to do work or to transfer heat.

  • Work is the force required to move an object against. It is the energy required to overcome friction and inertia.

  • Heat is the energy used to cause the temperature of an object to increase.

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Kinetic Energy

  • Kinetic energy is the energy of motion

    • Ek = ½ mv2

  • The kinetic energy of an object increases as the mass increases and as the square of the velocity increases.

  • A 1000 kg (2200 pounds) car traveling at 10 m/s (22.5 mph) has a kinetic energy of 50kJ. The same car traveling at 20 m/s (45 mph) has a kinetic energy of 200 KJ.

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    Potential Energy

    • Potential energy is the energy of position, it arises when a force operates on an object.

    • The two potential energies we are familiar with are gravitational and electrostatic.

    • Ep = mgh

    • Where m is the mass of the object, h is the height of the object relative to a reference height and g is the gravitational constant of 9.8 m/s2

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    Potential Energy (cont)

    • As chemists we don’t concern ourselves with the potential energy of gravitation due to the small size of the objects concerned.

    • Very concerned with electrostatic potential.

    • The electrostatic potential arises from the interactions between charged particles.

      • Eel = kQ1Q2/d

  • Where k is a constant of proportionality = 8.99 x 109 J-m/C2

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    Potential Energy (cont)

    • The chemical energy of molecules is due to the potential energy stored in the arrangements of their atoms.

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    Units of Energy

    • The SI unit of energy is the joule, J

    • The car from the previous example –

      • A 1000 kg (2200 pounds) car traveling at 10 m/s (22.5 mph)

      • Ek = ½ mv2

      • = ½ (1000 kg)(10 m/s)2

      • = ½ (1000 kg)(100 m2/s2)

      • = 50,000 kg-m2/s2

      • = 50 kJ

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    Units of Energy

    • The calorie was originally defined as the amount of energy required to raise the temperature of 1 g of water from 14.5 oC to 15.5 oC

    • 1 cal = 4.184 J (exactly)

    • Side note the calorie in food units is equal to 1 kcal.

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    System and Surroundings

    • We can not account for the entire universe when performing our analysis.

    • We focus on a small portion of the universe and that portion is called the system, everything else is the surroundings.

    • When we study the energy change that accompanies a chemical reaction, the reactants and products constitute the system.

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    • A closed system can exchange energy but not matter with the surroundings.

      • A car’s cylinder during ignition is a closed system.

    • An open system can exchange energy and matter with the surroundings.

      • A open beaker with boiling water

    • An isolated system can exchange neither energy nor matter with the surroundings.

      • An insulated thermos approximates an isolated system.

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    Transferring Energy: Work & Heat

    • Energy is transferred between systems and surroundings via heat or work or both.

    • Energy used to cause an object to move is called work.

      • W = F x d

    • We perform work when we lift an object against the force of gravity.

    • Heat is the energy transferred from a hotter object to a colder one. Heat always flows from hotter to colder objects.

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    The First Law of Thermodynamcis

    • Energy is conserved – energy can neither be created or destroyed only converted from one form to another.

    • The internal energy of a system is the sum of all the kinetic energy and potential energies of all its components.

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    Internal Energy

    • The internal energy includes translational energies, vibrational energies, rotational energies of every atom in the system.

    • The numerical value of the internal energy of a system is generally not known to a high degree of accuracy. (Why?)

    • We can calculate the change in internal energy.

      • ΔE = Efinal – Einitial

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    Internal Energy

    • ΔE contains three parts

      • A number

      • A unit

      • A direction

    • A negative value of ΔE indicates the system has lost energy to its surroundings

    • A positive value of ΔE indicates the system has gained energy from its surroundings

    • It is important to note that energy changes are from the point of view of the system not the surroundings.

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    Internal Energy (cont)

    • In a chemical reaction, the initial state of the system refers to the reactants.

    • The final state of the system refers to the products.

    • When hydrogen and oxygen form water at a given temperature, the system loses energy to the surroundings, thus the internal energy of the products is less than that of the reactants.

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    ΔE – heat and work

    • The internal energy of a system changes in magnitude as heat is added to or removed from the system or as work is done on it or by it.

    • ΔE = q + w

    • When heat is added to a system or work is done on a system, its internal energy increases.

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    Endothermic and Exothermic Processes

    • When a process occurs in which the system absorbs heat, the process is called endothermic.

    • During an endothermic process heat flows into the system from its surroundings

    • When a process occurs in which the system emits heat, the process is called exothermic.

    • During an exothermic process heat flows from the system to its surroundings.

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    State functions

    • The value of a state function depends only on the present state of the system, not on the path the system took to reach that state.

    • Because E is a state function, ΔE, depends only on the initial and final states of the system, not on how the change occurs.

    • E is a state function however q and w are not.

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    • Even though most of the chemical reactions we will examine occur at a constant atmospheric pressure, work is still being performed. (w = F x d)

    • As an example dissolving Zn in and acid release H2 gas. The expansion of the gas against atmospheric pressure is work.

    • The work involved in the expansion or compression of a gas is called pressure-volume work.

    • w = -P ΔV

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    Enthalpy (cont)

    • Enthalpy is a thermodynamic function that accounts for heat flow in a process occurring at constant pressure when the only work performed is P-V work.

    • Enthalpy is denoted by the symbol H

    • H = E + PV

    • The change in enthalpy of a system equals the heat gained or lost at constant pressure.

    • p 176

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    Enthalpies of reactions

    • Since ΔH = Hproducts – Hreactants the enthalpy change that accompanies a reaction is called the enthalpy of reaction or heat of reaction.

    • When exactly 2 moles of H2 reacts with exactly 1 mole of O2 to form 2 moles of H2O at a constant pressure, the system releases 483.6 kJ of heat.

    • 2H2 + O2 2H2O(g) ΔH = -483.6 kJ

    • Examine energy state diagram p 177

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    Thermochemical equations and enthalpy diagrams

    • Enthalpy is an extensive property.

      • The magnitude of ΔH is directly proportional to the amount of reactant consumed in the process.

      • CH4(g) + O2(g)  CO2(g) + 2H2O(l) ΔH = -890 kJ

      • If two moles of methane are consumed then

      • ΔH = -1780 kJ

    • The enthalpy change for a reaction is equal in magnitude, but opposite in sign, to the ΔH of the reverse reaction.

      • CO2 (g) + 2H2O(l)  CH4(g) + O2(g) ΔH = 890 kJ

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    Thermochemical equations and enthalpy diagrams (cont)

    • The enthalpy change for a reaction depends on the state of the reactants.

    • If the products in the combustion of methane were gaseous H2O instead of liquid H2O ΔH would be -802 kJ instead of – 890 kJ.

    • It is necessary to specify the state of the reactants and products.

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    • The value of ΔH can be determined experimentally via a calorimeter.

    • The measurement of heat flow is calorimetry.

    • We determine the magnitude of the heat flow by measuring the magnitude of the temperature change.

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    Statements of the Obvious

    • The more heat an object gains, the hotter it gets.

    • All substances change temperature when they are heated.

    • But less obvious is that the magnitude of the change with a given heat varies from substance to substance.

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    Heat Capacity and Specific Heat

    • The temperature change experienced by an object when it absorbs certain amount of heat is determined by its heat capacity, C.

    • The heat capacity of an object is the amount of heat required to raise its temperature by 1oC.

    • The heat capacity of a mole of a substance is it molar heat capacity. The heat capacity of one gram of a substance is its specific heat.

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    Specific Heat

    • Specific heat = (quantity of heat transferred)

    • (gram of substance)(temp. change)

      • Cs = q/(m x ΔT)

      • Exercise 5.5 p 181

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    • Constant-Pressure calorimetry is the process of measuring the heat transfer at constant, usually atmospheric, pressure.

    • Constant-Volume calorimetry is the process of measuring heat flow at constant volume in a device called a bomb calorimetery. (p 183)

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    Hess’s Law

    • Many enthalpies of reaction have been tabulated and from those tabulations it is possible to calculate the enthalpy of a reaction without the need for a calorimeter.

    • Hess’s law states that if a reaction is carried out in a series of steps, ΔH for the overall reaction will equal the sum of the enthalpies of reaction of the individual steps.

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    Enthalpies of Formation

    • Enthalpy of formation is the enthalpy or heat change during formation of a compound from its constituent elements.

    • The magnitude of any enthalpy change depends on the conditions of temperature, pressure and state.

    • Therefore to compare enthalpies we must define a set of conditions called the standard state.

    • STP – standard temperature and pressure

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    Standard enthalpy change

    • The standard enthalpy change of a reaction is defined as the enthalpy change when all reactants and products are in their standard states.

    • By definition, the standard enthalpy of formation of the most stable form of any element is zero.

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    • Example 5.10 pg 190

    • Using enthalpies of formation to calculate enthalpies of reaction

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    • The vast majority of the energy our bodies consume comes from carbohydrates and fats.

    • C6H12O6 + 6O2 6CO2 + 6H2O ΔH = -2803 kJ

    • Tristearin is a typical fat,

    • 2C57H110O6(s) + 163O2(g)  114 CO2 + 110 H2O(l) ΔH = -75,520 kJ

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    • Fossil Fuels

      • Natural gas

      • Petroleum

      • Coal