Energy

• Energy is the ability to do work or transfer heat (SI Unit: Joule)

• Thermodynamics is the study of energy and its transformations.

• Thermochemistry is the study of chemical reactions and the energy changes that involve heat.

• Potential energy - stored energy

• Kinetic energy - energy matter in motion

• Chemical energy is a form of potential energy.

First Law of Thermodynamics

• Energy is never created or destroyed, but it can be converted from one form to another.

• Chemical energy can be converted to heat to heat homes.

• Sunlight is converted to chemical energy in green plants.

Systems and Surroundings

• The portion of the universe that we single out to study is called the system.

• The surroundings are everything else.

• Example 1:

  • Molecules in a piston - the system.

  • The piston - the surroundings

• Example 2:

  • Ice - the system

  • Everything else - the surrounding

Types of Systems

• Open System: a region of the universe being studied that can exchange heat and mass with its surroundings.

• Closed System: a region of the universe being studied that can only exchange heat with its surroundings (not mass).

• Isolated System: a region of the universe that can not exchange heat or mass with its surroundings.

Internal Energy

• The internal energy of a system is the sum of all kinetic and potential energies of all components of the system; we use E to represent it.

• We don’t generally know E, only how it changes.

• The change in internal energy is \Delta E , the final energy of the system minus the initial energy of the system.

• Equation for change in energy: \Delta E=E_{final}-E_{initial}

  • Or \Delta E=E_{products}-E_{reactants}

Changes in Internal Energy

• If \Delta E>0, then E_{final} >E_{initial}

  • The system absorbed energy from the surroundings (endothermic)

• If \Delta E<0, then E_{final} <E_{initial}

  • The system released energy to the surroundings, which is an (exothermic)

• When energy is exchanged between the system and the surroundings, it is exchanged as either heat (q) or work (w).

  • Heat (q)

  • Work (w)

  • Equation for changes in internal energy: \Delta E=q+w

• The systems always results in an increase in the internal energy of a system when it gains heat and has work done on it by the surroundings

Thermodynamic Quantities

Thermodynamic quantities have three parts

• A number

• A unit

• A sign

• Note about the sign:

  • A positive (+) \Delta E results when the system gains energy from the surroundings.

    • Endothermic

  • A negative (-) \Delta E results when the system loses energy to the surroundings.

    • Exothermic

Change is energy, heat, work, and their signs

Exchange of Heat Between System and Surroundings

• When heat is absorbed by the system from the surroundings, the process is endothermic.

• When heat is released by the system into the surroundings, the process is exothermic.

State Functions

• We know that the internal energy of a system is independent of the path by which the system achieved that state.

• Internal energy is a state function.

• It depends only on the present state of the system, not on the path by which the system arrived at that state.

• \Delta E depends only on E initial and E final.

  • E is a state function

• q and w are not state functions

Work

• Usually the only work done by chemical or physical change is the mechanical work associated with a change in volume of gas.

• We can measure the work done by the gas if the reaction is done in a vessel that has been fitted with a piston: w=-P\Delta V

• The work is negative because it is work done by the system.

• Work = force x distance

Enthalpy

• Enthalpy (H) is the internal energy plus the product of pressure and volume (the sum of the energy change from a chemical reaction and the total work during the chemical reaction).

  • H=E+PV

• When the system changes at constant pressure, the change in enthalpy \Delta H,

  • \Delta H=\Delta E+P\Delta V

• A process is endothermic (absorbs energy) when \Delta H is positive.

• A process is exothermic (released energy) when \Delta H is negative.

Enthalpy of Reaction

• This quantity, \Delta H_rxn , is called the enthalpy of reaction (or heat of reaction).

The Truth About Enthalpy

• Enthalpy is an extensive property, meaning it does depend on the amount of material present.

• The enthalpy change for a reaction is equal in magnitude, but opposite in sign, to \Delta H for the reverse reaction.

  • A + B → C is \Delta H=+

  • C → A + B is \Delta H=-

• The enthalpy change for a reaction depends on the states of the reactants and the products (s,l, g, aq).

Calorimetry

• Calorimetry: the measurement of heat flow.

• Since we cannot know the exact enthalpy of the reactants and products, we measure \Delta H through calorimetry.

• Calorimeter: instrument used to measure heat flow.

• Calculating the amount of heat

  • q=mc\Delta T

    • q is heat

    • m is mass

    • c is the constant (specific heat)

      • Common constant is the specific heat of water (4.184\:J/g\cdot C)

    • \Delta T is change is temperature (final T - initial T)

  • Heat change between surroundings and the system

    • q_{soln }=-q_{rxn}

Hess’s Law

• We can calculate \Delta H using published \Delta H values and the properties of enthalpy.

• Hess’s law: If a reaction is carried out in a series of steps, \Delta H for the overall reaction equals the sum of the enthalpy changes for the individual steps.

• \Delta H is a state function, so for a particular set of reactants and products, \Delta H is the same whether the reaction takes place in one step or in a series of steps.

Enthalpies of Formation

• Enthalpy of formation (\Delta H_{f}): the enthalpy change for the reaction in which a compound is made from its constitute elements in their elemental forms.

Calculation of \Delta H

• We can use Hess’s law in this way: \Delta H=\sum n\Delta H_{f,products}-\sum m\Delta H_{f,reactan ts}^{°}, where n and m are the stoichiometric coefficients.

Bond Enthalpy

• The bond enthalpy is always positive because energy is required to break chemical bonds.

• Breaking bonds requires energy.

• Forming bonds releases energy.

• The greater the enthalpy, the stronger the bond.

Bond Enthalpies and Enthalpy of Reaction

• Add bond energy for all bonds made (+)

• Subtract bond energy for all bonds broken (-)

• We can predict whether a chemical reaction will be endothermic or exothermic using bond energies.

  • \Delta H_{rxn}=\sum bonds\:broken-\sum bonds\:formed

Reference: Chemistry The Central Science (14th Edition)