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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

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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.

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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.

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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.

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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.

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

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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

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K = °C + 273.15

°C = K – 273.15

°F = 9/5 °C + 32

°C = 5/9 (°F - 32)

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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

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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 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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Q = m x ∆t x cp

where Q = heat flowm = mass

∆t = change in temperaturecp = specific heat

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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

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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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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

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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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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

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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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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

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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.

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

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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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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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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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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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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

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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 – 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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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).

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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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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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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 = 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

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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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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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Energy is conserved only in elastic collisions

Momentum is conserved in every collision in which there is no friction.

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Gay-Lussac P ≈ T

Holding volume constant, the pressure is proportional to the absolute temp.

P1 / T1 = P2 / T2

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Boyle’s 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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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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Combined Gas Law:

P1V1/T1 = P2V2/T2

Ideal Gas Law:

PV = n R T (where n = # moles, and R = gas constant)

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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.

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

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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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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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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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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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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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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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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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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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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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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 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

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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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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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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.

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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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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

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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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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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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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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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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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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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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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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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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)

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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.

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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.

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