# The Kinetic Theory of Matter

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The Kinetic Theory of Matter
Chapter 10 The Kinetic Theory of Matter

10.1 Physical Behavior of Matter
States of Matter solid liquid gas

Intermolecular Forces (IMF)
Attractive forces between molecules. Much weaker than chemical bonds within molecules.

The Kinetic Theory of Matter
1. Matter is composed of PARTICLES. 2. Particle movement is rapid, constant, and random (Brownian motion)

The Kinetic Theory of Matter
3.All collisions are perfectly ELASTIC (NO energy lost).

Kinetic theory of matter
Kinetic energy (K.E.) = energy of motion gases have the least restriction on motion have the most K.E. solids have the most restriction on motion have the least K.E.

Kinetic model of gases Gases: matter with variable shape and variable volume Gas particles move in a straight line until they collide with container or each other

Kinetic model of gases Diffusion- random motion of gas particles that spread out and fill a space

Kinetic model of gases ideal gas—gases with perfectly elastic collisions and no intermolecular forces Most real gases behave as ideal gases except at very low temps and very high pressures

Kinetic model of gases pressure – force acting on a unit area of surface gas pressure— collisions of gas particles on objects

Kinetic model of gases atmospheric pressure— collisions of “air” particles on objects

STANDARD ATMOSPHERIC PRESSURE:
Kinetic model of gases SI unit of pressure = Pa (Pascal) standard pressure: (this is the “P” from STP) STANDARD ATMOSPHERIC PRESSURE: 1 atm = 760. mm Hg = 760. torr = kPa = 14.7 psi Psi = pounds per square inch mmHg = millimeter of mercury Atm – atmosphere kPa – kilopascals Torr - torr

examples of pressure conversions
1) Convert a pressure of 847 mm Hg to kPa. 847 mm Hg x kPa mm Hg 2) What is 8.9 psi expressed in atm? 8.9 psi x atm psi = 113 kPa = 0.61 atm

3) 344 mm Hg = _____ psi = 6.65 psi 344 mm Hg x 14.7 psi___

Kinetic model of liquids
Liquids: matter with variable shape and definite volume Particles slide past each other but are so close together they do not move in a straight line

Kinetic model of solids
solids: matter with definite shape and definite volume Particles cannot move past each other, they are in constant motion bouncing off neighbors

Crystalline solids Types of solids
a) crystal lattice—organized repeating pattern in 3-D b) unit cell—smallest repeating unit in a crystal

Crystalline solids continued
c) allotropes— two or more different arrangements for the same element in the same state (C: graphite, diamond)

Types of solids 2.amorphous— solids without a set structure
incomplete crystal lattice formed b) rubber, plastics, glass Candles, peanut butter, cotton candy

Other forms of matter 1. amorphous solids 2. liquid crystals—an intermediate phase formed when solids partially melt in only one or two dimensions and can flow like a liquid (LCD = liquid crystal display)

Other forms of matter 3. Plasmas
gaseous mixture of ions -exists at high temperatures most common form of matter in the universe but least common on Earth itself

Plasmas continued an ionized gas that conducts electricity -forms at very high temps when matter absorbs energy and breaks apart The sun is made of plasma - also found in fluorescent lights

10.2 – Kinetic energy and changes of state
Temperature and kinetic energy temperature—the measure of the average K.E. of particles in a sample Kelvin (K) – SI base unit of temperature; measures average K.E.

Temperature and kinetic energy
When temp increases, particle motion increases. When temp decreases, particle motion decreases. A temp of 300 K has twice the kinetic energy as 150 K.

Temperature and kinetic energy
0 Kelvin = absolute zero = no molecular motion No degrees sign ( ° ) is used with Kelvin numbers There will never be negative numbers for Kelvin temperatures!.

Temperature Scales Fahrenheit Celsius Kelvin Anders Celsius 1701-1744
Lord Kelvin (William Thomson)

Temperature Scales Notice that 1 kelvin = 1 degree Celsius Fahrenheit
Boiling point of water 32 ˚F 212 ˚F 180˚F 100 ˚C 0 ˚C 100˚C 373 K 273 K 100 K Freezing point of water Notice that 1 kelvin = 1 degree Celsius

Element Freezing Point, ºC Boiling Point, ºC Oxygen -219 -183 Chlorine -101 -34 Nickel 1455 2913 Phosphorus 44 280 1. Which of the elements are gases at 50ºC? At -50ºC? 2. Which of the elements are liquids at 50ºC? At -50ºC? 3. Which of the elements are solids at 50ºC? At -50ºC? 4. Which element has the smallest temperature range as a liquid? The largest temperature range?

Converting Temperature
Kelvin-Celsius conversion equation K = C Express K in degrees Celsius. K = C = C C = 93 °C Convert a temperature of 45 °C to Kelvin. K = C K = = 318 K

Fahrenheit / Celsius Formulas
°F = 9/5 °C (°F - 32) * 5/9 = °C A person with hypothermia has a body temperature of 29.1°C. What is the body temperature in °F? °F = 9/5 (29.1°C) = = 84.4°F

Changing states vaporization - conversion of a liquid to a gas or vapor below the boiling point (b.p.) evaporation rate – depends on surface area, temp, and humidity

b) condensation—conversion from a gas or vapor to a liquid

Changing states c) sublimation—changing from a solid directly to a vapor (w/o becoming liquid first) EX: dry ice, mothballs, solid air fresheners

d) deposition—changing from a vapor/gas directly to a solid (w/o becoming liquid first)

What type of phase change is occuring?
1. solid carbon dioxide(dry ice)to carbon dioxide gas 2. ice to liquid water 3. liquid bromine to bromine vapor 4. liquid water to ice 5. water vapor to liquid water sublimation melting vaporization freezing condensation

Vapor Pressure and boiling
Vapor Pressure - pressure of vapor above a liquid at equilibrium high vapor pressure = volatile volatile = easily evaporates The greater the fraction of molecules which can escape the liquid, the greater the vapor pressure

At some point in time the number of vapor molecules rejoining the water equals the number leaving to go into the vapor phase

Vapor pressure and boiling point
Boiling Point - temp at which v.p. of liquid equals external pressure -depends on atmospheric pressure & IMF Normal B.P. - b.p. at 1 atm When the vapor pressure of a liquid equals atmospheric pressure, the liquid has reached its boiling point, which is 100°C for water at sea level. At this point, molecules throughout the liquid have the energy to enter the gas or vapor phase.

Effects of Intermolecular Forces (IMF)
When IMF’s are weak vapor pressure is high volatility is high boiling point is low

Heat of Vaporization Joule (J) – the SI unit of energy
heat of vaporization – the amount of heat necessary to boil 1 mole of a substance at its boiling point

Heat of Fusion Melting point – temp of a solid when it becomes a liquid = freezing point (temp when liquid becomes a solid) heat of fusion – energy needed for 1kg of a substance to solidify at it’s freezing point

Freezing/Melting point
B. Heating Curves Gas Boiling point Liquid Freezing/Melting point Solid

Changing state IMPORTANT: temp does not change during a phase change.
Increasing the temp will only make the change happen faster.

Phase Diagrams Shows the phases of a substance at different temps and pressures.

triple point -the point on a phase diagram that represents the temperature and pressure at which three phases of a substance can coexist. All six phase changes can occur at the triple point: freezing and melting, evaporation and condensation, sublimation and deposition.

Phase Diagrams critical point -the critical pressure and critical temperature above which a substance cannot exist as a liquid.

THE END!