Presentation on theme: "Chapter 10 The Kinetic Theory of Matter. 10.1 Physical Behavior of Matter States of Matter – solid – liquid – gas."— Presentation transcript:
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 gasgases 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 = kPa = 14.7 psi
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 1 atm psi = 113 kPa = 0.61 atm
3) 344 mm Hg = _____ psi 344 mm Hg x 14.7 psi___ 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 latticeorganized repeating pattern in 3-D b) unit cellsmallest 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 a)incomplete crystal lattice formed b) rubber, plastics, glass Candles, peanut butter, cotton candy
Other forms of matter 1. amorphous solids 2. liquid crystalsan 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 temperaturethe 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 FahrenheitFahrenheit CelsiusCelsius KelvinKelvin Anders Celsius Lord Kelvin (William Thomson)
Temperature Scales 1 kelvin = 1 degree Celsius Notice that 1 kelvin = 1 degree Celsius Boiling point of water Freezing point of water Celsius 100 ˚C 0 ˚C 100˚C Kelvin 373 K 273 K 100 K Fahrenheit 32 ˚F 212 ˚F 180˚F
ElementFreezing Point, ºCBoiling Point, ºC Oxygen Chlorine Nickel Phosphorus Which of the elements are gases at 50 º C? At -50 º C? 2. Which of the elements are liquids at 50 º C? At -50 º C? 3. Which of the elements are solids at 50 º C? At -50 º C? 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 + 32 (°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) + 32 = = = 84.4°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) condensationconversion from a gas or vapor to a liquid
Changing states c) sublimationchanging from a solid directly to a vapor (w/o becoming liquid first) EX: dry ice, mothballs, solid air fresheners
d) depositionchanging 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.sublimation 2.melting 3.vaporization 4.freezing 5.condensation
Vapor Pressure and boiling Vapor Pressure 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
Effects of Intermolecular Forces (IMF) When IMFs are weak – vapor pressure is high 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 its 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.