1) The Nature of Energy (KE, PE, heat, and work)

• Energy

  • allows us to do work or produce heat

  • heat transfer from chemical processes

  • JOULES

    

    amount needed to accelerate

  • Law of Conservation

    • energy can be converted from one form to another but can be neither creater nor destroyed

  • total energy content of the universe is constant

• Types of energy

  • Potential energy

    • energy due to position or composition (chemical bonds)

    • can result from attractive and repulsive forces

    • stored energy

  • Kinetic energy

    • energy due to the motion of an object (atoms or molecules)

    • depends on the mass of the object (m) and its velocity (v)

    • KE = ½ mv²

• Conversion of energy

  • transfering energy through physical contact

• Methods of transfering energy

  • heat

    • transfer of energy between two objects due to a temperature difference (hot to cold)

    • temperature reflects random motion of particles in a substance (kinetic energy) - C and K

    • energy absorbed = increase temperature

    • energy released = decrease temperature

  • work

    • force acting over a distance

    • work = force x distance = F x ^h

• Parts of the universe

  • universe = system + surroundings

  • system

    • part of the universe on which one wishes to focus their attention

    • in the lab (where you are)

    • reactants and products of a reaction in their container

    • 3 types

      • open system

        • open system can exchange mass and energy, usually in the form of heat with its surroundings

      • closed system

        • which allows transfer of energy (heat) not mass

      • isolated system

        • does not allow the transfer of either mass or energy

  • surroundings

    • include everything else in the universe

    • things other than the reactants and products in their container

• Types of Reactions

  • Exothermic

    • results in the evolution of heat (energy)

    • energy flows out of the system

  • Endothermic

    • results in the absorption of energy from the surroundings

    • heat flows into a system

• Reaction mechanism

  • energy gained by the surroudings must be equal to the energy lost by the system

  • endothermic reactions result from a lowered potential energy of the reaction system

    • must add energy from surroundings

    • PE products > PE reactants

  • exothermuc reactions, potential energy stored in chemical bonds is converted to thermal energy via heat

    • energy must be released to surroundings

    • PE products < PE reactants

• change (triangle) PE

  • stored in the bonds of products as compared with the bonds of reactants

  • exo - more energy is released while forming new bonds than is consumed while breaking the bonds in the reactants

    

  • endo - energy that flows into the system as the heat is used to increase the potential energy of the system (energy needed to break bonds)

    

• thermodynamics

  • study of energy and its interconverstions

  • 1st law of thermodynamics

    • energy of the universe is constant (law of conservation of energy)

• Internal energy (E) of a system

  • sum of kinetic and potential energies of all particles in a system

  • can be changed by flow of work, heat, or both

    

• Parts of thermodynamic quantities

  • number

    • gives the magnitude of chnage

  • sign

    • indicates the direction of flow (endo/exo)

    • reflects the systems point of view

  • in endo

    • q = +x

    • when the surroundings do work on the system

    • w is positive

  • in exo

    • q = -x

    • when a system does work on surroundings

    • w is negative

      

• Work

  • associated with a chemical process

  • work done by a gas through expansion

  • work done to a gas through compression

  • equation

    • F=force

    • P= pressure (force per unit of area)

    • P= F/A

    • H= height

    • A = area

• Gas expansion

  • delta V positive

  • W is negative

• Gas compression

  • delta V is negative

  • W positive

2) Enthalpy and Calorimetry

• Enthalpy

  • delta H- the heat content of a system

    • a state function that is defined as

    • H=E+PV

    • E= internal energy of system

    • P= pressure of the system

    • V= volume of the system

  • enthalpy is a state function because it does not depend on the pathway between states

  • at constant pressure enthalpy change (delta H = q_p )

    • q_p = heat at constant pressure

    • delta H = enthalpy of products - enthaply of reactant

      • positive = endo

      • negative = exo

• Calorimetry

  • science of measuring heat (q)

    • based on observations of temperature change when a body absorbs or discharges energy in form of heat

  • calorimeter

    • device used to determine the heat associated with a chemical reaction

  • heat capacity ( C )

    • heat absorbed/increase in temperature

  • specific heat capacity

    • energy required to raise the temperature of one gram of a substance by one degree celisius

      • j/c*g or j/k*g

  • molar heat capacity

    • energy required to raise the temperature of one mole of a substances by one degree celsius

    • j/c*mol or j/k*mol

    • for metals different from water - less energy

      

• Constant - pressure calorimetry

  • atmospheric pressure remains constant during the process

  • used to determine enthalpy changes in reactions in a solution

  • delta H = q_p

• calculation of heat (q) for a neutralization reaction: released= absorbed

  • specific heat capacity x mass of solution x change in temperature

  • s x m x delta T

  • heat of reaction is an extensive property - depends on amount tho

• Constant - volume calorimetry

  • used in conditions when experiments are to be performed under constant volume

  • no work done since V must chnage for PV to work to be performed

    • bomb calorimeter

  • delta E = q + w = q = q_v

3) Hess’s Law

• Enthalpy change as a state function

  • going from reactants to products the ehthalpy (delta H) is the same whether the reaction takes place in one step or in a series of steps

• If the reaction is reversed the enthalpy is also reverse (endo and exo)

• magnitude of enthalpy is directly proportional to the quantities of reactants and products in a reaction

  • if the coefficients in a balanced reaction are multiplied by an interger the value of enthalpy is multiplied by the same integer

  • somehow combine all of the the equations and combine the enthalpy values

    • watch out for coefficients

4) Standard Enthalpies of Formation

• Delta H_f^o

  

• change in the enthalpy that accompanies of one mole of a compound from its elements with all substances in their standard states

  • standard state

    • precisely defined reference states

    • Gas state

      • 1 atm

    • Pure substance Condensed state (l or s)

      • pure liquid or solid

    • in Solution

      • 1M

    • Standard state of an element

      • form where element exists under conditions of 1 atm and 25C

      • Oxygen = standard state O₂(g)

      • ^H in any lone element is 0

        

  • degree symbol on thermodynamic function shows that the corresponding process that is carried out under standard conditions

  • elements in its standard state is 0

5) Sources of Energy

• w

• e

• e

• e

• ee