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
2
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.
3
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.
4
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.
5
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.
6
Thermal equilibrium is the situation in which no heat moves from one object to another (they have the same temperature).
7
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
8
Fahrenheit Celsius KelvinB
oiling Pt. H2O 212 100 373B
ody temp 98.6 37 310F
reezing 32 0 273C
oincidence -40 -40 233A
bsolute zero -460 -273 0
10
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
11
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
12
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
14
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).
15
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
16
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.
17
Metal Specific Heat
Thermal Conductivity
Density Electrical Conductivity
cp (cal/g° C)
k (watt/cm K)
(g/cm3) 1 E 6/ Ώm
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
19
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
20
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
21
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.
22
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.
23
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
24
Matter is defined as any material that has mass, occupies volume, and exhibits inertia (resistance to movement).
25
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.
26
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
27
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
28
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
29
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
30
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
31
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
32
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).
33
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.
34
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)
35
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
36
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
37
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.
38
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)
39
Energy is conserved only in elastic collisions
Momentum is conserved in every collision in which there is no friction.
40
Gay-Lussac P ≈ T
Holding volume constant, the pressure is proportional to the absolute temp.
P1 / T1 = P2 / T2
41
Boyle’s Law V ≈ 1/P
If the temp. is held constant, the volume of a gas varies inversely with the pressure
P1V1 = P2V2
42
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
43
Combined Gas Law:
P1V1/T1 = P2V2/T2
Ideal Gas Law:
PV = n R T (where n = # moles, and R = gas constant)
44
Chemical properties are those properties of a substance that can be determined by a chemical test. They are seen by the material’s tendency to change, either alone or by interaction, and in doing so form different materials.
45
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
46
Physical properties are those properties used in identifying substances when we use our senses. These do not require chemical analysis.
47
Color, hardness, density, texture, magnetic attraction, solubility, taste, light transmission, viscosity, refractive index, specific heat, boiling point, melting-freezing point, odor, expansion-contraction coefficients
48
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.
49
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
50
The molecular makeup (specific arrangement of atoms) is changed, resulting in new substances being formed and energy changes occurring.
Two types: exothermic and endothermic
51
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
52
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
53
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
54
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
55
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
56
Dry ice sublimates
CO2 + H2O + sunlight glucose
Air in heated tire expands
Burning coal
Water frozen into ice
Acid dissolves metal
57
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),paradichlorobenzene, camphor, iodine crystals, CO2 fire extinguishers
58
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.
59
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.
60
Boiling point is defined as the temperature at which the liquid’s 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.
61
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).
62
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
65
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.
66
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.
67
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
68
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)
69
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.
70
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.
71
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
72
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
73
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)
74
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.
75
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.
76