Thermochemistry

1
Thermochemistry
2
Heat Temperature

• Joseph Black explained heat in terms of a fluid
  (Lavoisier had called this fluid caloric from
  Latin word for heat.
• Count Rumford friction could convert mechanical
  energy into heat (motion as cause)
• John Dalton idea of atoms

3
Heat History

• James Prescott Joule tried to find the
  mechanical equivalent of heat (where a given
  amount of energy produces the same amount of
  heat)
• James Clerk Maxwell developed a solid
  explanation showing relationship between motion
  of atoms and heat.

4
Heat

• Heat flows from hot to colder areas due to a
  temperature difference only (till thermal
  equilibrium is established).
• Heat is a form of internal energy which is
  transferred from one object to another due to a
  difference in temperature between the objects.

5
Heat Content

• The heat content of a substance is the total
  energy of all the particles of that substance.
• The total energy combines both kinetic and
  potential energies.

6
Temperature

• The temperature of a body of matter is a measure
  of the average kinetic energy of the random
  motion of its particles.
• Temperature is that property of a substance which
  determines whether it is in thermal equilibrium
  with another object.

7
Thermal Equilibrium

• Thermal equilibrium is the situation in which no
  heat moves from one object to another (they have
  the same temperature).

8
Thermometers

• Thermometers work on idea of thermal expansion
  the amount of expansion or contraction is always
  the same for the same increase or decrease in
  temperature.
• 3 types gas (air), liquid (Hg alcohol), solid
  (bimetallic)
• Know creation and calibration ideas

9
Temperature Conversions

• K C 273.15
• C K 273.15
• F 9/5 C 32
• C 5/9 (F - 32)

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Scale Comparisons

• Fahrenheit Celsius Kelvin
• Boiling Pt. H2O 212 100
  373
• Body temp 98.6 37
  310
• Freezing 32 0
  273
• Coincidence -40 -40
  233
• Absolute zero -460 -273 0

11
Heat Units

• 15 calories the amount of heat needed to
  raise 1 gram of water from 14.5 to 15.5 C at
  1 atmosphere of pressure
• kilocalorie kcal or Calorie 1000 cal
• 1 calorie 4.185 Joule
• 1 kcal 4185 Joule

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Specific Heat

• Specific heat is the amount of heat needed to
  raise 1 gram of water 1 C at 1 atmosphere of
  pressure
• What is the degree change if 1 calorie of heat is
  added to 1 gram samples of
• water helium ice gold

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Heat Flow (Q)

• Q m x ?t x cp
• where Q heat flow m mass
• ?t change in temperature cp specific
  heat

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Principle of Heat Exchange

• The amount of heat lost by a substance is equal
  to the amount of heat gained by the substance to
  which it is transferred.
• m x ?t x cp m x ?t x cp
• heat lost heat gained

15
Specific Heat Notes

• Specific heat how well a substance resist
  changing its temperature when it absorbs or
  releases heat
• Water has high cp results in coastal areas
  having milder climate than inland areas (coastal
  water temp. is quite stable which is favorable
  for marine life).

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More Specific Heat

• Organisms are primarily water thus are able to
  resist more changes in their own temperature than
  if they were made of a liquid with a lower cp

17
Water and Heat

• When calories of heat are added to water there is
  a small change in temperature because most of the
  heat energy is used to disrupt hydrogen bonds
  before water molecules can begin to move faster.
• Temp. of water drops many additional hydrogen
  bonds form, releasing a considerable amount of
  heat energy.

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Absolute Zero
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Material Data
Metal Specific Heat Thermal Conductivity Density Electrical Conductivity
cp (cal/g C) k (watt/cm K) (g/cm3) 1 E 6/ ?m
Brass 0.09 1.09 8.5
Iron 0.11 0.803 7.87 11.2
Nickel 0.106 0.905 8.9 14.6
Copper 0.093 3.98 8.95 60.7
Aluminum 0.217 2.37 2.7 37.7
Lead 0.0305 0.352 11.2
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Heat Transfer Mechanisms

• Conduction faster vibrating particles collide
  with less energetic neighbor and transfer energy
  to it
• Convection motion of hot fluid, displacing cold
  fluid in path setting up convection current
• Radiation energy transmitted by electromagnetic
  waves

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Thermal Expansion of Water

• From 0 to 4 the volume of water in a sample
  decreases (the greatest density is at 4 c)
• Ice floats body of water freezes from top down
  allowing life underneath to continue

22
Ice Open Structure

• Water mlcl can participate in 4 bonds with other
  water mlcl (solid mlcl can have as many as a
  dozen bonds with surrounding mlcl resulting in a
  more compact substance).
• The spaces between mlcl in ice are greater than
  the same spaces in liquids.

23
Density of Ice

• Density of ice increases from 0 to 4 as large
  clusters of mlcl break into smaller clusters that
  takes up less space in the aggregate. Above 4
  normal thermal expansion is seen with a decrease
  in density.

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Latent Heat

• Heat of Fusion amount of heat needed to change
  solid to liquid at its melting point
• Heat of Vaporization heat needed to change
  liquid to gas at boiling point
• Heat of Sublimation heat to change a solid to
  gas
• Heat of Condensation heat released when gas
  condenses to a liquid

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Matter

• Matter is defined as any material that has mass,
  occupies volume, and exhibits inertia (resistance
  to movement).

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States of Matter

• Solids definite shape and volume, resist
  deformation
• Very close spacing of particles
• Particles appear to vibrate around fixed points
• Particles vibrate faster at higher temp.

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Types of Solids

• Crystalline particles arranged in regular,
  repeated patterns (long-range order) example
  NaCl (s)
• Amorphous solids that lack the definite
  arrangement of crystalline solids (have
  short-range order)
• Examples pitch, glass, plastics

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Liquids

• Definite volume, resist compression, take shape
  of container
• Greater spacing between particles, particles
  appear to travel in straight line paths between
  collisions but appear to rotate and/or vibrate
  about moving points

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Gases

• Have no definite shape or volume, take shape and
  volume of container
• Can be compressed or dispersed, particles vibrate
  very rapidly, relatively far apart
• There are no intermolecular forces holding
  particles together

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Plasma

• Very high temperature ionized gas
• No fixed volume or shape
• Most are mixtures that are not easily containable
• Particles are electrically charged and of low
  density
• Example the Milky Way

31
Energy

• Energy having the ability to do work
• Work a push or pull over a distance
• Force a push or pull
• Momentum mass x velocity
• Linear momentum of a moving body is a measure of
  its tendency to continue in motion at a constant
  velocity

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Potential Energy

• Potential Energy the energy a body possesses by
  virtue of its position, composition, and/or
  condition
• P.E. is the stored energy
• P.E. mass x gravity x height

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Kinetic Energy

• K.E. the energy of motion
• K.E. is conserved in all elastic collisions
• K.E. ½ m v2 (m mass, v velocity)
• Heat energy flows from hot objects to cooler ones
  by transfer of K.E. when particles collide
  (conduction).

34
Intermolecular Forces

• P.E. forces that hold mlcl together and in
  correct position in solids.
• P.E. forces that hold mlcl together in liquids.
• These forces are between mlcl.
• Gases have enough K.E. to prevent formation of
  these forces.

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Kinetic Molecular Theory of Gases

• Gases are mlcl in continuous motion.
• An increase in temp. increases speed thus
  increasing K.E.
• All gases are compressible
• Gases display diffusion
• Gases can be liquified (called liquifaction)

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Closed System Info ? Pressure, Volume, Temp.

• Nothing escapes or enters system
• All mlcl in motion (have K.E.)
• Mlcl exert uniform pressure against walls of
  container
• Mlcl exert pressure on other mlcl as they
  collide, push, bounce off other mlcl

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Pressure

• Pressure Force / Area
• Atmospheric Pressure cumulative net force per
  area generated by weight of our atmosphere
• Values 14.7 lb/in2, 101.3 kPa,
• 760 mm of Hg, 1 atm, 1033 g/cm2

38
Gas Pressure

• The pressure a gas exerts on the walls of its
  container is the sum of the forces acting ( the
  frequency of collisions plus the force of each
  collision) due to the random collisions of near
  limitless numbers of moving molecules.

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Collisions

• Inelastic collisions the normal type in which
  objects lose energy and slow down
• Elastic collisions particles bounce off,
  exchange energies but there is no loss of energy
  (energy is conserved but may be redistributed)

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Conservation in Collisions

• Energy is conserved only in elastic collisions
• Momentum is conserved in every collision in which
  there is no friction.

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Gas Laws

• Gay-Lussac P T
• Holding volume constant, the pressure is
  proportional to the absolute temp.
• P1 / T1 P2 / T2

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Gas Laws

• Boyles Law V 1/P
• If the temp. is held constant, the volume of a
  gas varies inversely with the pressure
• P1V1 P2V2

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Gas Laws

• Charles Law V T
• If the pressure is held constant, the volume of a
  gas is proportional to its absolute temp.
• V1 / T1 V2 / T2
• For every degree increase in temp. the volume
  increases by 1/273 of its original volume

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Better Gas Law Equations

• Combined Gas Law
• P1V1/T1 P2V2/T2
• Ideal Gas Law
• PV n R T (where n moles, and R
  gas constant)

45
Chemical Properties

• Chemical properties are those properties of a
  substance that can be determined by a chemical
  test. They are seen by the materials tendency
  to change, either alone or by interaction, and in
  doing so form different materials.

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Chemical Properties

• Does the substance support combustion? Burn
  itself?
• How does it react with acids? With oxygen?
  With electricity?
• Examples alcohol burns, wood decays, sodium
  explodes and burns in water

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Physical Properties

• Physical properties are those properties used in
  identifying substances when we use our senses.
  These do not require chemical analysis.

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Physical Properties

• Color, hardness, density, texture, magnetic
  attraction, solubility, taste, light
  transmission, viscosity, refractive index,
  specific heat, boiling point, melting-freezing
  point, odor, expansion-contraction coefficients

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Physical Changes in State

• This is a change in the physical properties of a
  substance without a change in the chemical
  composition.
• The arrangement of molecules may be changed but
  the molecular makeup remains the same.
• These changes involve intermolecular forces which
  increase or decrease during the change.

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Physical Changes

• Ice (0 C) heat ? steam (100 C)
• 36 g 25 920 cal 36 g
• Steam (100 C) ? ice (0 C) heat
• 36 g 36 g 25 920 cal

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Chemical Changes in State

• The molecular makeup (specific arrangement of
  atoms) is changed, resulting in new substances
  being formed and energy changes occurring.
• Two types exothermic and endothermic

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Exothermic Chemical Changes

• Any chemical change that releases energy
• The amount released must be greater than the
  amount used to start reaction
• Bond making is exothermic (energy is released
  into surroundings

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Exothermic Examples

• Oxidation wooden splint burning (giving off
  light, heat, CO2, H2O
• Burning H2 in air, body reactions, dissolving
  metals in acid, mixing acid and water, sugar
  dehydration, plaster of Paris in water

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Endothermic Reactions

• Any chemical change that absorbs energy
• Energy continues to be absorbed as long as
  reaction continues
• Bond breaking is endothermic (energy is absorbed
  from surroundings

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Endothermic Examples

• Electrolysis (breaking water down into H2 and O2
  by running electricity in it)
• Photosynthesis, pasteurization, canning
  vegetables
• 2 H2 O2 ? 2H2O energy
• 4 g 32 g 36 g 136 600 cal
• 2H2O energy ? 2H2 O2
• 36 g 136 600 cal 4g 32 g

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Changes using Energy

• Physical change strength of intermolecular
  forces increased or decreased
• Chemical change bonds formed or broken
• Energy absorbed bonds broken or intermolecular
  forces overcome
• Energy released bonds formed or intermolecular
  forces strengthened

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Examples

• Dry ice sublimates
• CO2 H2O sunlight ? glucose
• Air in heated tire expands
• Burning coal
• Water frozen into ice
• Acid dissolves metal

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Sublimation

• Sublimation is the direct change of a solid to a
  gas
• Deposition is the change of a gas to a solid
• Examples moth balls (naphthalein),paradichlorobe
  nzene, camphor, iodine crystals, CO2 fire
  extinguishers

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Liquid-Solid Phase Change

• Melting-freezing point this is the same
  temperature at which a pure substance can change
  into solid from liquid or solid into liquid.
• The solid-liquid phases are in equilibrium.

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Melting-Freezing Point

• When heat is added to a solid, the temp. will
  increase till it reaches the melting-freezing
  point. It will remain at that temp. until all
  the solid has melted and then the temp. can rise
  again according to its specific heat.

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Liquid Gas Phase Change

• Boiling point is defined as the temperature at
  which the liquids vapor pressure is equal to
  outside (atmospheric pressure usually).
• When the vapor pressure equal atmospheric
  pressure as many mlcl are leaving the surface as
  are re-entering the surface of the liquid.

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Boiling Point

• Boiling point varies with elevation.
• Cooking times must adjust due to elevation.
• Pressure cookers can cook food more rapidly due
  to increased pressure, resulting in high boiling
  points (which cooks food faster).

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Heating Curve
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Phase Change Diagram
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Thermochemistry In-Depth

• Study of heat changes that take place in a change
  of state or chemical reaction
• If heat is released, process is called exothermic
• If heat is absorbed, process is called endothermic

66
Heat

• Heat is energy transferred from one object to
  another due to a difference in temperature.
• We measure the temperature change that
  accompanies heat transfer.
• We have to measure the temperature change of the
  surroundings (the solvent, container,
  atmosphere).
• The system is the reactants and products of the
  reaction.

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Situations

• When a system releases heat to surroundings, the
  temperature of the surroundings increases
  (exothermic). An example would be combustion of
  propane in a barbecue grill.
• When a system absorbs heat, the temperature of
  the surroundings decreases (endothermic). An
  example would be melting ice.

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Enthalpy

• The amount of heat transferred depends on the
  energy stored in each substance. This stored
  energy is called heat content or enthalpy and is
  represented by H.
• ?H qp Enthalpy heat transferred
• qp m ?t cp

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Specific Heat

• cp reflects that ability of a substance to absorb
  heat (defined as the amount of heat needed to
  raise the temperature of 1 gram by 1 degree
  Celsius)
• cp of water 1.00 cal/g C or 4.185 J/g C
• In most situations it is the temperature change
  of the surroundings that is measured (which
  equals the heat releases/absorbed from the
  reaction itself)

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Exothermic Reactions

• For the increase in temperature of the
  surroundings, heat must be released by the
  system.
• The surroundings increase is positive while the
  heat release by the system must be negative.
• Exothermic reactions always have negative values.

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Endothermic Reactions

• Heat absorbed by system results in temperature
  decrease for surroundings (negative quantity).
• Heat absorbed by system must have positive value.
• Enthalpy change for endothermic is always a
  positive value.

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Heat of Solution

• Heats of solution deal with the process of a
  solute dissolving in a solvent.
• In the case of an ionic solute, there are two
  processes
• Energy to break apart the ionic bonds in the
  crystal lattice (called crystal lattice energy)
• Energy released when the free ions form
  attractive forces with water molecules (called
  heat of hydration)
• The heat of solution is the sum of these two
  effects

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Heat of Solution Example

• Crystal lattice energy of KCl
• KCl (s) ? K1 (g) Cl1- (g) ?H
  167.6 kcal
• Heat of hydration of KCl
• K1 (g) Cl1- (g) ? K1 (aq) Cl1- (aq)
• 
  ?H - 163.5 kcal
• Overall KCl (s) ? K1 (aq) Cl1- (aq)
• ?H 4.1 kcal - Endothermic
  Reaction
• 
  

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Heats of Solution (kcal/mol)

• NH4NO3 6.1 (endothermic)
• NaOH - 10.6 (exothermic)
• KNO3 8.0 (endothermic)
• KClO3 9.89 (endothermic)
• KOH - 13.77 (exothermic)
• NaCl 0.93 (endothermic)
• NaC2H3O2 4.085 (endothermic)

75
Neutralization Reactions

• Usually these reactions are exothermic but adding
  vinegar to baking soda is slightly endothermic.
• The neutralization reaction is slightly
  exothermic.
• HC2H3O2 (aq) NaHCO3 (aq) ? CO2 (g)
  NaC2H3O2 (aq) H2O (l)
• net bond formation
• Evaporation of the liquid occurs as the CO2
  escapes from solution. Evaporation absorbs heat,
  cooling the liquid (along with expansion of
  bubbles also helps to cool the surroundings)
  net result is endothermic reaction.

76
Addition of Acid to Water

• Mixing a strong acid with water is exothermic.
• Breaking chemical bonds requires energy.
• Forming chemical bonds releases energy.
• HCl (g) ? H1 (aq) Cl1- (aq)
• It looks like heat would be absorbed because the
  bond between the H and Cl is broken. The
  hydrogen reacts with water to form a complex
  H3O(H2O) n (where n is between 1 and 9).
• This hydration makes the overall reaction
  strongly exothermic.