chapter 10 the kinetic theory of matter. 10.1 physical behavior of matter states of matter – solid...

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

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 = 101.3 kPa = 14.7 psi

examples of pressure conversions

1) Convert a pressure of 847 mm Hg to kPa.

847 mm Hg x 101.3 kPa 1 760. mm Hg

2) What is 8.9 psi expressed in atm?

8.9 psi x 1 atm 1 14.7 psi

= 113 kPa

= 0.61 atm

3) 344 mm Hg = _____ psi

344 mm Hg x 14.7 psi___ 1 760. mm Hg

= 6.65 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

Types of solids

1) Crystalline 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 solids2.amorphous— solids without a set

structurea)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 matter3. 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 ScalesTemperature Scales

• FahrenheitFahrenheit

• CelsiusCelsius

• KelvinKelvin

Anders Celsius1701-1744

Lord Kelvin(William Thomson)1824-1907

Temperature ScalesTemperature Scales

Notice that 1 kelvin = 1 degree Celsius1 kelvin = 1 degree Celsius

Boiling point of Boiling point of waterwater

Freezing point of Freezing point of waterwater

CelsiusCelsius

100 ˚C100 ˚C

0 ˚C0 ˚C

100˚C100˚C

KelvinKelvin

373 K373 K

273 K273 K

100 K100 K

FahrenheitFahrenheit

32 ˚F32 ˚F

212 ˚F212 ˚F

180˚F180˚F

Element Freezing Point, ºC

Boiling Point, ºC

Oxygen -219 -183

Chlorine -101 -34

Nickel 1455 2913

Phosphorus 44 2801. 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 + 273

Express 366.13 K in degrees Celsius.K = C + 273 366.13 = C + 273 C = 93

°C

Convert a temperature of 45 °C to Kelvin.K = C + 273 K = 45 + 273 = 318 K

Fahrenheit / Celsius FormulasFahrenheit / Celsius Formulas

°F = 9/5 °C + 32 °F = 9/5 °C + 32

(°F - 32) * 5/9 = °C(°F - 32) * 5/9 = °C

A person with hypothermia has a body temperature of A person with hypothermia has a body temperature of 29.1°C. What is the body temperature in °F? 29.1°C. What is the body temperature in °F?

°F °F = = 9/5 (29.1°C) + 329/5 (29.1°C) + 32

= = 52.4 + 3252.4 + 32

= = 84.4°F84.4°F

Changing states

a)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 statesc) 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)

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

What type of phase change is occuring?

1.sublimation2.melting3.vaporization4.freezing5.condensatio

n

Vapor Pressure and boiling

• Vapor PressureVapor 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

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

B. Heating Curves

Freezing/Melting point

Solid

Liquid

Boiling point

Gas

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!