Standard Enthalpies of Formation

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Standard Enthalpies of Formation

• The standard enthalpy of formation of a
  substance, denoted DHfo, is the enthalpy change
  for the formation of one mole of a substance in
  its standard state from its component elements in
  their standard state.

• Note that the standard enthalpy of formation for
  a pure element in its standard state is zero.

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Standard Enthalpies of Formation

• The law of summation of heats of formation states
  that the enthalpy of a reaction is equal to the
  total formation energy of the products minus that
  of the reactants.

• S is the mathematical symbol meaning the sum
  of, and m and n are the coefficients of the
  substances in the chemical equation.

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A Problem to Consider

• You record the values of DHfo under the formulas
  in the equation, multiplying them by the
  coefficients in the equation.

• You can calculate DHo by subtracting the values
  for the reactants from the values for the
  products.

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How is the heat of sublimation, ?Hsub, the
enthalpy change for the reaction H2O(s) ?
H2O(g) related to ?Hfis and ?Hvap?
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Fuels

• Food fills three needs of the body

• It supplies substances for the growth and repair
  of tissue.
• It supplies substances for the synthesis of
  compounds used in the regulation of body
  processes.
• It supplies energy. About 80 of the energy we
  need is for heat. The rest is used for muscular
  action and other body processes

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Fuels

• A typical carbohydrate food, glucose (C6H12O6)
  undergoes combustion according to the following
  equation.

• One gram of glucose yields 15.6 kJ (3.73 kcal)
  when burned.

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Fuels

• A representative fat is glyceryl trimyristate,
  C45H86O6. The equation for its combustion is

• One gram of fat yields 38.5 kJ (9.20 kcal) when
  burned. Note that fat contains more than twice
  the fuel per gram than carbohydrates contain.

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Figure 6.15 Sources of energy consumed in the
United States (1996).
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Fuels

• Fossil fuels account for nearly 90 of the energy
  usage in the United States.

• Anthracite, or hard coal, the oldest variety of
  coal, contains about 80 carbon.
• Bituminous coal, a younger variety of coal,
  contains 45 to 65 carbon.
• Fuel values of coal are measured in BTUs (British
  Thermal Units).
• A typical value for coal is 13,200 BTU/lb.
• 1 BTU 1054 kJ

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Fuels

• Natural gas and petroleum account for nearly
  three-quarters of the fossil fuels consumed per
  year.

• Purified natural gas is primarily methane, CH4,
  but also contains small quantities of ethane,
  C2H6, propane, C3H8, and butane, C4H10.
• We would expect the fuel value of natural gas to
  be close to that for the combustion of methane.

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Fuels

• Petroleum is a very complicated mixture of
  compounds.

• Gasoline, obtained from petroleum, contains many
  different hydrocarbons, one of which is octane,
  C8H18.

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Fuels

• With supplies of petroleum estimated to be 80
  depleted by the year 2030, the gasification of
  coal has become a possible alternative.

• First, coal is converted to carbon monoxide using
  steam.

• The carbon monoxide can then be used to produce a
  variety of other fuels, such as methane.

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Practice Problem 6.45
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Practice Problem 6.46
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States of Matter Liquids and Solids

• Chapter 11

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Exam Friday will cover Chapters 5 and 6 and as
much of 11 as we finish tomorrow.
Suggested Problems Chapter 11 21, 23, 29, 34,
37, 39, 43, 45, 49, 51, 53, 55, 59, 61, 63, 67,
69, 71, 95, 97, 119
Online homework for Ch 5 and 6 due today
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Figure 11.1 Dry ice.Photo courtesy of American
Color.
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States of Matter

• Comparison of gases, liquids, and solids. (see
  Figure 11.12)

• Gases are compressible fluids. Their molecules
  are widely separated.
• Liquids are relatively incompressible fluids.
  Their molecules are more tightly packed.
• Solids are nearly incompressible and rigid. Their
  molecules or ions are in close contact and do not
  move.

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Figure 11.2 Representation of the states of
matter.
Lattice energy
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Changes of State

• A change of state or phase transition is a change
  of a substance from one state to another. (see
  Table 11.1)

gas
liquid
solid
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Figure 11.3 A beaker containing iodine
crystals. Photo courtesy of James Scherer.
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Figure 11.3 Iodine has appreciable vapor
pressure below its melting point. Photo courtesy
of James Scherer.
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Vapor Pressure

• Liquids are continuously vaporizing.

• If a liquid is in a closed vessel with space
  above it, a partial pressure of the vapor state
  builds up in this space.
• The vapor pressure of a liquid is the partial
  pressure of the vapor over the liquid, measured
  at equilibrium at a given temperature. (see
  Figure 11.4)

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Figure 11.4 Measurement of the vapor pressure
of water.
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Figure 11.5 Distribution of kinetic energies of
molecules in a liquid.
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Figure 11.6 Rates of vaporization and
condensation of a liquid over time.
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Vapor Pressure

• The vapor pressure of a liquid depends on its
  temperature. (see Figure 11.7)

• As the temperature increases, the kinetic energy
  of the molecular motion becomes greater, and
  vapor pressure increases.
• Liquids and solids with relatively high vapor
  pressures at normal temperatures are said to be
  volatile.

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Figure 11.7 Variation of vapor pressure with
temperature.
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Boiling Point

• The temperature at which the vapor pressure of a
  liquid equals the pressure exerted on the liquid
  is called the boiling point.

• As the temperature of a liquid increases, the
  vapor pressure increases until it reaches
  atmospheric pressure.
• At this point, stable bubbles of vapor form
  within the liquid. This is called boiling.
• The normal boiling point is the boiling point at
  1 atm.

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Figure 11.8 Boiling of a liquid.
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Freezing Point

• The temperature at which a pure liquid changes to
  a crystalline solid, or freezes, is called the
  freezing point.

• The melting point is identical to the freezing
  point and is defined as the temperature at which
  a solid becomes a liquid.
• Unlike boiling points, melting points are
  affected significantly by only large pressure
  changes.

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Heat of Phase Transition

• To melt a pure substance at its melting point
  requires an extra boost of energy to overcome
  lattice energies.

• The heat needed to melt 1 mol of a pure substance
  is called the heat of fusion and denoted DHfus.

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Figure 11.9 Heating curve for water.
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Heat of Phase Transition

• To boil a pure substance at its melting point
  requires an extra boost of energy to overcome
  intermolecular forces.

• The heat needed to boil 1 mol of a pure substance
  is called the heat of vaporization and denoted
  DHvap. (see Figure 11.9)

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A Problem to Consider

• The heat of vaporization of ammonia is 23.4
  kJ/mol. How much heat is required to vaporize
  1.00 kg of ammonia?

• First, we must determine the number of moles of
  ammonia in 1.00 kg (1000 g).

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A Problem to Consider

• The heat of vaporization of ammonia is 23.4
  kJ/mol. How much heat is required to vaporize
  1.00 kg of ammonia?

• Then we can determine the heat required for
  vaporization.

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Figure 11.11 Phase diagram for water (not to
scale).
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Phase Diagrams

• A phase diagram is a graphical way to summarize
  the conditions under which the different states
  of a substance are stable.

• The diagram is divided into three areas
  representing each state of the substance.
• The curves separating each area represent the
  boundaries of phase changes.

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Phase Diagrams

• Below is a typical phase diagram. It consists of
  three curves that divide the diagram into regions
  labeled solid, liquid, and gas.

.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
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Phase Diagrams

• Curve AB, dividing the solid region from the
  liquid region, represents the conditions under
  which the solid and liquid are in equilibrium.

.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
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Phase Diagrams

• Usually, the melting point is only slightly
  affected by pressure. For this reason, the
  melting point curve, AB, is nearly vertical.

.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
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Phase Diagrams

• Curve AC, which divides the liquid region from
  the gaseous region, represents the boiling points
  of the liquid for various pressures.

.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
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Phase Diagrams

• Curve AD, which divides the solid region from the
  gaseous region, represents the vapor pressures of
  the solid at various temperatures.

.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
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Phase Diagrams

• The curves intersect at A, the triple point,
  which is the temperature and pressure where three
  phases of a substance exist in equilibrium.

.
B
C
solid
liquid
pressure
.
gas
A
D
temperature
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Phase Diagrams

• The temperature above which the liquid state of a
  substance no longer exists regardless of pressure
  is called the critical temperature.

.
B
C
solid
liquid
pressure
.
gas
A
D
Tcrit
temperature
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Phase Diagrams

• The vapor pressure at the critical temperature is
  called the critical pressure. Note that curve AC
  ends at the critical point, C.

.
B
Pcrit
C
solid
liquid
(see Figure 11.13)
pressure
.
gas
A
D
Tcrit
temperature
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Figure 11.13 Observing the critical phenomenon.
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Figure 11.12 Phase diagrams for carbon dioxide
and sulfur (not to scale).
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Properties of Liquids Surface Tension and
Viscosity

• The molecular structure of a substance defines
  the intermolecular forces holding it together.

• Many physical properties of substances are
  attributed to their intermolecular forces.
• These properties include vapor pressure and
  boiling point.
• Two additional properties shown in Table 11.3 are
  surface tension and viscosity.

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Figure 11.18 A steel pin floating on the surface
of water.
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Figure 11.19 Liquid levels in capillaries.
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Figure 11.20Comparison of the viscosities of
two liquids. Photo courtesy of James Scherer.
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Properties of Liquids Surface Tension and
Viscosity

• Surface tension is the energy required to
  increase the surface area of a liquid by a unit
  amount.

• This explains why falling raindrops are nearly
  spherical, minimizing surface area.
• In comparisons of substances, as intermolecular
  forces between molecules increase, the apparent
  surface tension also increases.