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Thermochemistry! PowerPoint PPT Presentation

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Thermochemistry!. AP Chapter 5. Temperature vs. Heat. Temperature is the average kinetic energy of the particles in a substance. Heat is the energy that is transferred from one object to another. Heat always flows from the hotter object to the colder object. Energy!.

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AP Chapter 5

Temperature vs. Heat

  • Temperature is the average kinetic energy of the particles in a substance.

  • Heat is the energy that is transferred from one object to another.

  • Heat always flows from the hotter object to the colder object.


  • Energy is the ability to do work.

  • Kinetic Energy- the energy of motion

  • Potential Energy – the energy that an object has as a result of its composition or its position with respect to another object.

Units of Energy

  • 1 Joule = 1 kg m2/s2 (1 kJ is 1000J)

    • Used to calculate the energies associated with chemical reactions.

  • Calorie – Amount of energy required to raise the temperature of 1 gram of substance 1 °C. (This is specific heat!)

    1 calorie will raise the temperature of 1 g of H2O from 14.5 °C to 15.5 °C.

    1 calorie is equal to 4.184 Joules (exactly!)

  • Systems and Surroundings

    • System – the portion used in a study.

      • It can be an open system or a closed system.

    • Opensystem – matter and energy can interact with the surroundings.

    • Closedsystem – the matter cannot interact with the surroundings.

    First Law of Thermodynamics

    • Energy Is Conserved!

    Internal Energy

    • Internal Energy is the sum of all the kinetic and potential energies of all its components.

      ΔE = Efinal - Einitial


    • A positive value for ΔE is when Efinal > Einitial

    • If energy has been absorbed from its surroundings, it is endothermic.

    • If energy is given off to the surroundings, it is exothermic.

    Initial state refers to the reactants, while final state refers to the products.

    Endothermic reaction

    Exothermic reaction

    A system composed of H2 (g) and O2 (g) has greater internal energy than a system composed of H2O (l).

    Gases have greater kinetic energy and must lose some of that energy to change states back to the liquid state.

    Internal energy is a function of state.

    • If a battery is shorted out and loses energy to the environment only as heat, no work is done.

    • If a battery is discharged and loses energy as work (to make the fan run) it also loses heat energy.

    • The value of ∆E is the same.


    • The change in enthalpy for a reaction (∆H) is the overall measure of energy that is absorbed to break bonds and the energy that is released when new bonds form.

    • A reaction is said to be spontaneous if it occurs without being driven by an outside force. (driving forces are enthalpy & entropy)

    • ∆H = ΣH(products) - ΣH(reactants)

    In an endothermic system where it absorbs heat, ∆H will be positive (∆H > 0).

    In an exothermic system, where heat is given off, ∆H will be negative (∆H < 0).

    Enthalpy Diagrams

    • Enthalpy is an extensive property – it depends on how much you have. If 1mol of CH4 and 2 mol O2 yield -890 kJ, then 2 mol CH4 and 4 mol O2 would yield double that.

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


    • This is a measure of the amount of energy that is needed or lost when a certain mass of a substance changes temperature.

    • q = mC∆T

    • q is the amount of energy (J)

    • m is the mass of the substance (g)

    • C is the specific heat capacity of the substance

    • ∆T is the change in temperature


    • Calorimeters are devices that measure the transfer of heat from one object to another.

    Heat of Formation (∆H°f)

    • The heat change that occurs when one mole of a compound is formed from its elements at 1 atm pressure.

    • Generally, the standard enthalpy of formation for any element in its most stable form is 0. (i.e. O2 gas would have a standard enthalpy of 0.)

    • Remember Appendix C!

    Standard Enthalpy Changes

    • The standard enthalpy change can be calculated from the standard enthalpies of formation of the reactants and products in the reaction (see Appendix C for values.)

    • ∆H°rxn = Σn∆H°f(products) - Σm ∆H°f(reactants)

    • The n and m refer to the molar coefficients in the chemical equation.

    Also refer to Appendix C!

    Hess’s Law

    • If you can break a chemical reaction into several steps, add up all of the ∆H’s for each step to get the overall ∆H for the reaction.


    • Entropy is a measure of randomness or disorder of a system. The greater the disorder, the greater the entropy.

    • In terms of entropy, gases>liquids>solids.

    • When pure substance dissolves in a liquid, its entropy increases.

    • When gas molecules escape a solvent, entropy increases.

    • Entropy increases with molecular complexity.

    • Reactions that increase the number of moles of particles often increase the entropy of the system.


    • Na+(aq) + Cl-(aq) → NaCl (s)∆S is negative

    • NH4Cl (s)→ NH3(g) + HCl (g) ∆S is positive

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