Chemistry and Energy

1
Chemistry and Energy
2
What is Energy?

• Standard definitionAbility to accomplish a task,
  the capacity to do work
• Work is applying a force through a distance
• Thusly, work/energy can only be measured when
  they are transferred, moving from one place to
  another

3
Kinetic Energy

• The energy of moving objects
• Large objects, such as cars, baseballs, people,
  we can see the motion
• Small objects, such as the molecules and atoms
  of chemistry, we cannot see, but their motion is
  what causes the temperature of the larger objects
• Even in air, the motion of particles is making
  itself known in the temperature
• Matter is in constant random motion

4
Kinetic Energy

• Calculation in physics is done with KE1/2 mv2
• where mmass and vvelocity.
• Both have their effects
• If a baseball hits you, it hurts more if a major
  league pitcher throws it than a 3rd grader
• If a large objects hits your car at 5 mph, it
  will cause much more damage than a 3rd grader on
  a bicycle

5
Kinetic Energy

• In chemistry, we relate KE to Temperature,
  KE3/2 RT, where Rthe gas constant and TKelvin
  temperature
• Temperature is the measure of the average kinetic
  energy of particles. In a sample, some
  individual molecules may be well above or below
  the average
• The higher the average is, the more energy can be
  transferred, which means it could cook something
  faster or cause a greater burn to your skin

6
Potential Energy

• Energy of position or statewhere it is or what
  it is
• Position has to do with gravityrock on top of a
  hill or water behind a dam
• State has to do with energy stored in some waya
  bowstring or a springor a chemical bond
• In chemistry, we are interested in the potential
  energy of bonds which can be released, in food,
  in explosives, in fuel

7
Transfer of Energy

• Occurs by way of forms of Kinetic Energy
• By heat energyit flows from areas of high
  temperature to areas of low temperature.
• The units used depend on the ones using it
• Engineers and researchers use the
• joulewhich is the SI unit
• Biologists and biochemists use the calorie,
  which is the energy required to raise the
  temperature of 1 gram of water by 1oC

8
Transfer of Energy

• Measuring heat transfer requires knowing how well
  the materials involved will absorb heat. This is
  called the Specific Heat, and it tells how much
  energy is needed to raise the temperature by 1oC.
  For water, this value is 4.18 Joules/ gram oC,
  which is very high compared to most other
  substances
• The formula for heat transfer is qmcpDT
• where qheat, mmass, cpspecific heat value,
  and DTchange in temperature

9
Measurement of Energy

• By using energy transferred from substances to
  water, we can measure how much energy they
  contain
• Food substances are measured with a bomb
  calorimeter, which actually explodes the food
  in an all oxygen atmosphere
• Food calories are actually 1000 regular
  calories, or kilocalories
• Calorimeters used in chemistry classes are of the
  coffee cup type, where a measured amount of
  water is used to measure the heat exchange

10
Coffee Cup Calorimeter
Bomb Calorimeter
11
Energy for Your Body

• Your body needs a constant source of energy,
  averaging about 1500 kcal/day minimum just to
  keep you alive
• Any excess energy you consume in food over that
  amount goes to fuel your activity or to be stored
  as fat.
• Any type of food can become fat, just if there is
  enough extra Calories, 3500 from any source
  becomes one pound of fat.

12
Specific Heat Capacity

• For any substance, the amount of heat necessary
  to raise the temperature of that substance by 1
  C
• heat capacity per gram J/C g

13
Example

• If a sample of iron with a mass of 10.0 g changes
  from 50.4 C to 25.0 C when 114 J of heat are
  released, what is its specific heat capacity?

14
Homework 16-a

• p. 495 4-6, 8, 11
• p.524ff 46, 79, 80

15
System and Surroundings

• System That on which we focus attention
• Surroundings Everything else in the universe
• Universe System Surroundings

16
Exo and Endothermic

• Heat exchange accompanies chemical reactions.
• Exothermic Heat flows out of the system (to
  the surroundings). Source is from bond energies
  differences between reactants and products.
• Endothermic Heat flows into the system (from
  the surroundings).

17
Enthalpy Considerations

• DH reaction H products - H reactants
• Enthalpy is an extensive property, the greater
  the number of moles of substances, the greater
  the enthalpy

18
Calorimetry

• Calorimetry is the method of measuring heat flow
  from any system by measuring the difference in
  temperature of an absorbing substance, usually
  water. The value of q
• from calorimetry is the opposite of DH
  calculated from reaction enthalpies.
• q - DH (mH2O )(4.18 J/g oC)(DT)
• or
• q - DH (H.C.bomb)(DT) for bomb calorimeter

19
State Changes

• Energy changes cause state changes, from solid to
  liquid to gas and reverse
• A process which absorbs energy is called
  endothermic. This happens as the substances
  molecules get more energetic
• A process where energy is released is called
  exothermic. When energy is released, the
  molecules slow down and are attracted to each
  other to become liquid or solid

20
Melting Point

• Molecules break loose from lattice points and
  solid changes to liquid. (Temperature is
  constant as melting occurs.) The reverse is
  freezing
• vapor pressure of solid vapor pressure of
  liquid

21
Boiling Point

• Constant temperature when added energy is used to
  vaporize the liquid. The reverse process is
  condensation.
• vapor pressure of liquid pressure of
• surrounding atmosphere

22
(No Transcript)
23
Conservation of Energy

• First Law of ThermodynamicsEnergy cannot be
  created or destroyed, only converted between
  formsThe total energy of the Universe is
  conserved
• Energy is converted from concentrated useful
  forms into less useful forms (light and heat)
  which spread out too far to be useful
• Work processes are very inefficientcars use only
  20 of the energy in the fuel for moving. Even
  the human body is only 44 efficient, most of the
  energy becomes heat or fat.

24
Spontaneous Processes and Entropy

• Thermodynamics lets us predict whether a process
  will occur but gives no information about the
  amount of time required for the process.
• A spontaneous process is one that occurs without
  outside intervention. No indication of rate is
  involved.

25
Entropy

• Second Law of ThermodynamicsQuality of the
  energy available is continually
  degradingbecoming less useful as it spreads.
  This spreading is called ENTROPYa measure of the
  randomness or disorder of a system
• Entropy is always increasing in the universe and
  the tendency is to increase in any system unless
  energy is input to reverse it. In your body,
  parts and molecules to wear out, so you must
  input food to provide new materials and energy to
  rebuild. The energy in food comes ultimately
  from the SUN, which is continually losing useable
  energy out into space.

26
Positional Entropy

• Entropy brings about an increase in possible
  positions. A gas expands into a vacuum because
  the expanded state has the highest positional
  probability of states available to the system.
• Therefore,
• Ssolid lt Sliquid ltSaqueousltlt Sgas

27
The Second Law of Thermodynamics

• . . . in any spontaneous process there is always
  an increase in the entropy of the universe.
• ?Suniv gt 0
• for a spontaneous process.
• ?Suniv ?Ssystem ?Ssurroundings
• ?Ssystem can decrease if ?Ssurroundings increases
  more

28
Gibbs Free Energy

• Gibbs free energy (DG) is the total energy of any
  process
• DG relates all aspects of kinetic energy, the
  rotational, vibrational, and translational, as
  well as entropy.

29
Free Energy

• ?G ?H ? T?S (from the standpoint of the
  system)
• A process (at constant T, P) is spontaneous in
  the direction in which free energy decreases
• ??G means ?Suniv

30
Effect of ?H and ?S on Spontaneity
31
Homework 16-b

• p. 519 39, 40, 41, 44
• p. 524ff 51, 53, 54, 55, 63, 66, 69, 74, 90

32
Chemical Kinetics

• The area of chemistry that concerns reaction
  rates.

33
Reaction Rate

• Change in concentration (conc) of a reactant or
  product per unit time.

Instantaneous Rate--slope of the curve at any one
point
34
Collision Model

• Key Idea Molecules must collide to react.
• However, only a small fraction of collisions
  produces a reaction. Why?
• Arrhenius An activation energy must be overcome.

35
Arrhenius Statement

• Collisions must have enough energy to produce the
  reaction (must equal or exceed the activation
  energy).
• Orientation of reactants must allow formation of
  new bonds.

36
Exothermic
Endothermic
37
Factors Which Affect Rates

• Nature of the Reactants
• Surface area of Solids
• Concentration of Reactants
• Temperature of Reaction Mixture
• Catalysts

38
(No Transcript)
39
(No Transcript)
40
Catalysis

• Catalyst A substance that speeds up a reaction
  without being consumed
• Enzyme A large molecule (usually a protein)
  that catalyzes biological reactions.
• Homogeneous catalyst Present in the same phase
  as the reacting molecules.
• Heterogeneous catalyst Present in a different
  phase than the reacting molecules.

41
(No Transcript)
42
Homework 17

• p. 535 4, 6, 7
• p. 541 11, 12, 13
• p. 554 ff 35, 37, 41, 44, 45, 47, 64

43
Chemical Equilibrium

• The state where the concentrations of all
  reactants and products remain constant with time.
• On the molecular level, there is frantic
  activity. Equilibrium is not static, but is a
  highly dynamic situation.

44
(No Transcript)
45
Reversible Reactions

• Reactions may not go totally to completion in the
  forward direction, but seem to stop and reverse
  direction, depending on their spontaneity in the
  forward direction. Such a reaction is said to be
  reversible, at least to a minor amount. When the
  forward and reverse rates are identical, chemical
  equilibrium has been reached

46
The Law of Mass Action

• For any reaction, such as
• jA kB ? lC mD
• The law of mass action represents the balance
  between reactants and products by the equilibrium
  expression

47
Equilibrium Expression Example

• 4NH3(g) 7O2(g) ? 4NO2(g) 6H2O(g)

48
Notes on Equilibrium Expressions (EE)

• The Equilibrium Expression for a reaction is the
  reciprocal of that for the reaction written in
  reverse.
• When the equation for a reaction is multiplied by
  n, EEnew (EEoriginal)n
• The units for K depend on the reaction being
  considered.
• The value of K determines which side of the
  reaction is going to be greater Kltlt1, reactants
• K gtgt 1 products
• K close to 1, balanced

49
Heterogeneous Equilibria

• . . . are equilibria that involve more than one
  phase.
• CaCO3(s) ? CaO(s) CO2(g)
• K CO2
• The position of a heterogeneous equilibrium does
  not depend on the amounts of pure solids or
  liquids present.

50
Homework 18-a

• p. 567 2 p. 568 6, 7
• p. 590ff 27, 30, 33, 52, 53

51
Le Châteliers Principle

• . . . if a change is imposed on a system at
  equilibrium, the position of the equilibrium will
  shift in a direction that tends to reduce that
  change.

52
Effects of Changes on the System

• 1. Concentration The system will shift away
  from the added component. The system will shift
  toward from the removed component.
• 2. Temperature K will change depending upon the
  temperature (treat the energy change as a
  reactant).

53
Effects of Changes on the System (continued)

• 3. Pressure
• a. Addition of inert gas does not affect the
  equilibrium position.
• b. Decreasing the volume shifts the
  equilibrium toward the side with fewer moles.

54
(No Transcript)
55
Summary of LeChatelier
Concentration Equilibruim shifts away from what
is added, toward what is removed. Temperature
Equilibrium shifts away from added temperature,
depending on type of reaction, Exo-toward
reactants Endo-toward products Pressure
Equilibrium shifts toward less molecules when
pressure is added
56
Example 13.14

• Arsenic can be extracted from its ores by first
  reacting the ore with oxygen to form solid As4O6,
  which is then reduced with carbon
• As4O6 (s) 6 C (s) As4 (g) 6 CO (g)
• Predict the effect of these changes upon the
  position of equilibrium
• Adding carbon monoxide
• Adding or removing C or As4O6
• Removing gaseous As4
• LeChateliers principle states that addition
  shifts away from what is added, so (a) shifts
  toward the reactants, that is left. Solids being
  added have no effect, so (b) remains the same.
  Removing any component shifts toward what is
  removed, so (c) shifts right.

57
Example 13.15 16

• 15Predict the shift in equilibrium when volume
  is reduced in each of these(causing increased
  pressure)
• P4 (s) 6 Cl2 (g) 4 PCl3 (l)
• Since there are no moles of gas on the right,
  volume reduction will favor that side.
• b. PCl3 (g) Cl2 (g) PCl5 (g)
• Since there are fewer moles of gas on the right,
  the equilibrium will shift that direction.
• c. PCl3 (g) 3 NH3 (g) P(NH2)3 (g) 3 HCl (g)
• Since there are 4 moles on either side of the
  equation, there is no shift due to pressure
  change.
• 16For each of these reactions, predict the shift
  of equilibrium with an increase in temperature
• N2 (g) O2 (g) 2NO (g) DHº 181 kJ
• This reaction is endothermic, so heat acts as
  an added reactant, favoring the products to the
  right
• 2 SO2 (g) O2 (g) 2 SO3 (g)
  DHº -198 kJ
• This reaction is exothermic, heat a product, so
  it shifts to the left.

58
Homework 18-b
p. 574 11, 12, 13, 14, 15 p. 590ff 37, 41, 44,
46, 59-64