The Kinetic Theory of Matter explains the properties of solids, liquids, & gases
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The Kinetic Theory of Matter explains the properties of solids, liquids, & gases. The Kinetic Theory of Matter. Based on idea that particles of matter are in constant motion. Describes properties of solids, liquids, & gases in terms of the FORCE of the particles

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The Kinetic Theory of Matter explains the properties of solids, liquids, & gases

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The Kinetic Theory of Matter explains the properties of solids, liquids, & gases

The Kinetic Theory of Matter

Based on idea that particles of matter are in constant motion.

Describes properties of solids, liquids, & gases in terms of the FORCE of the particles

The constant random motion of tiny particles called Brownian motion

Physical Behavior of Matter

Section 10.1

States of Matter

  • Four states of matter


  • Particles-closely packed; can’t be compressed

  • Voids - extremely small

  • Particle motion is vibratorymotion; definite shape & volume

  • Apply heat-particles vibrate more & move SLIGHTLY farther apart; causes solid to expand

Kinetic Model of Solids

STRONG intermolecular forces result in rigid structure of solids

Particles move (vibrate) but not past each other

Particles occupy fixed 3-D positions that repeat throughput the solid

3-D arrangement = crystal lattice


  • Flowing matter w/ definite volume & indefinite shape

  • Particles have weak bonds that keep them close; more space to move; particles able to move relative to each other

  • When heat applied, liquids expand a little

Kinetic Model of Liquids

Particles of liquid slide past each other; consider magnetized spheres

Intermolecular forces maintain their volume NOT shape


  • Flow, too (consider wind)

  • Particles far apart; complete freedom of movement

  • Motion is random

  • No definite shape/volume

  • Easily compressed into smaller volume

  • Expand & contract in response to temperature changes more so than liquids & solids

Kinetic Model of Gases

Particles in gas in constant, random motion

Change direction ONLY when they strike wall of container OR another gas particle (Air-hockey puck)

Density (M/V) in gas lower than solid

Few particles in gas vs solid of same volume (due to space between particles)

5 Assumptions of the Kinetic Theory

Gases are made of molecules in constant, random movement.

LARGE portion of the volume of a gas = empty space. The volume of all gas molecules, in comparison, is negligible.

5 Assumptions of the Kinetic Theory

The molecules show no forces of attraction or

repulsion (UNLIKE solids & liquids).

No energy is lost in collision of molecules; the impacts are completely elastic.

The temperature of a gas is the average KE of all of

the molecules.

Ideal Gases

  • Ideal gases = gases that obey the assumptions of the kinetic theory

  • Except for temperatures extremes, most real gases behave like ideal gases

  • At temperature extremes, forces between particles & particle size begin to matter

  • At temperature extremes gases no longer follow the assumptions of the kinetic theory

Gases & Pressure, Temperature, & Volume

  • KToM explains gas pressure = the total force exerted by gas molecules colliding against the walls of a container.

  • IF the container can expand, like a balloon/tire,

    in pressure can the volume; THUS the balloon/tire will get BIGGER .

  • If you the temperature of the gas, the KE of its molecules &, the pressure/volume

Gases & Pressure, Temperature, & Volume

  • There is a relationship between pressure, volume and temperature in an ideal gas

  • If you the pressure & hold the volume constant, the temperature (principle of a refrigerator)

  • If you the temperature & hold the pressure constant, the volume (heating a balloon)

Earth’s Atmosphere and Pressure

  • We’re at the bottom of an ocean of air

  • Atmospheric pressure = force exerted on us by molecules of air (14.7 lbs/square inch)

  • Atmospheric pressure related to column of air;

  • As elevation ↑ pressure↓

  • As elevation ↓, pressure ↑

  • How does atm pressure affect YOU?

Other Forms of Matter

  • Amorphous Solids = arrangement of molecules is fairly random; so, crystal lattice is loosely packed ; haphazard, disjointed

  • Examples = GLASS,

cotton candy


Liquid Crystals

  • When solids melt, the crystal lattice disintegrates; particles lose their 3-D pattern

  • Liquid crystals-NOT liquid OR solid

  • When melted LCs lose their rigid organization in 1 or 2 dimensions NOT all 3 dimensions

  • Interparticle forces in liquid crystals are relatively weak; when forces in lattice are broken, crystals can flow like liquids

  • Liquid crystal displays (LCDs) used in TVs, watches, calculators, thermometers, etc.


  • Form at very temperatures

  • Plasma = gas that has been energized; some e- break free from, but travel w/their nucleus

  • Plasma = free e- & ions of that element.

  • Gases can become plasmas in several ways, ALL include pumping the gas w/ energy.

  • Examples = stars, fluorescent tubes, neon lights, etc.

Examples of Plasma

Energy and Changes of State

Section 10.2

Temperature & Kinetic Energy & Particle Motion

  • Temperature = measure of the average kinetic energy of particles in a material

  • When heated liquid & gas particles have more kinetic energy BUT not all particles have the same kinetic energy; particles are moving at different speeds

  • Generally, as temperature matter moves to a more active state; as temperature matter moves to a less active state

The Kelvin Scale

  • Absolute zero = temperature at which a substance would have zero (or very little) kinetic energy

  • Kelvin Scale = used for temperature; it is defined so temperature of a substance is directly proportional to the average kinetic energy of the particles

  • 0 on Kelvin scale = absolute zero & measured as Kelvins; divisions on Fahrenheit & Celsius scale are measured in degrees

  • Celsius degree & Kelvins = the same size; absolute zero = -273.150C

  • Kelvin scale measures everything ABOVE absolute zero; all numbers are positive

Temperature Conversions

  • When converting from kelvin (K or TK) to Celsius (C or TC), and vice versa, the magic number is 273!!

  • K= (0C+ 273); K= (150C+ 273) = 288 K

  • 0C= (K- 273); 0C= (320 K - 273) = 470C

Mass & Speed of Particles

KE of gas depends on mass & speed of particles

1. Gases at SAME temp have SAME average KE

2. LARGER gas molecule simply moves SLOWER than SMALL gas molecule

Ex: O2 = 16x more massive than H2; at SAME temp, H2 moves FASTER than O2

Mass & Speed of Matter

  • Random motion causes particles to spread out to fill a container

  • DIFFUSION = the process in which these particles fill a space

  • Particles move from areas of high concentration to areas of low concentration

  • Rate of diffusion of a gas dependent upon the KE of that gas/substance

Changes of State

TRIPLE POINT = The single specific temperature & pressure at which all 3 phases can co-exist

CRITICAL POINT = The conditions where gas & liquid become indistinguishable

Different phases of a system may be represented using a phase diagram.

Axes of the diagrams are typically pressure & temperature

Phase Diagram for Water

Changes of State


  • particles of a liquid form a gas by escaping from the surface

  • 3 things affect evaporation rate? Area of the surface, temperature, humidity

  • Volatile liquids evaporate quickly (perfumes, paint)

  • As liquids evaporate, they cool

Heating Curve-based on standard temp & pressure

Changes of State


The process of heating a SOLID substance to a point where it turns LIQUID.

FREEZING is the opposite of melting. It is the process of REMOVING heat from a liquid & turning the liquid into a solid.

The freezing point is the SAMETEMPERATURE as the melting point.

Heating Curve-based on standard temp & pressure

Changes of State

  • Sublimation-process by which particles in a solid change to gaseous state w/o melting

  • Condensation-reverse of evaporation; gaseous particles become close (condense) & form a liquid

Specific Heat

  • To change the temperature of a SOLID

    2.1 Joules/g0C

  • To change the temperature of a LIQUID

    4.2 Joules/g0C

  • To change the temperature of a GAS

    2.02 Joules/G0C

    Heat = mass x specific heat x temperature change

    Q = m x c x (Tf –Ti)

Heat of Vaporization

  • The amount of heat required (absorbed by the liquid) to convert unit mass of a liquid into its vapor w/o a change in temperature.

  • 2260 Joules = ENREGY needed to move the molecules in 1 g of water FAR enough apart that they form water vapor (JOULE, J,= SI unit of energy required to lift a 1-g mass 1m against the force of gravity)

  • Heat of Vaporization (Hv) of H2O = 2260 J/g

Heat of Vaporization of Water Hv = 2260 J /g

  • The diagram right shows the uptake of heat by 1 kg of H2O, from ice at -50 ºC  to steam above 100 ºC.

    A: Rise in temp. as ice absorbs heat.B: Absorption of latent heat of fusion.C: Rise in temp. as liquid H2O absorbs heat.D: Water boils & absorbs latent heat of vaporization.E: Steam absorbs heat & thus increases its temperature.

Heat of Fusion

  • The heat nrg which must be removed to solidify a liquid or added to melt a solid

  • Melting point=temperature of the solid when its crystal lattice begins to break apart (intermolecular forces are overcome & solid becomes a liquid)

  • Freezing Point= temperature of liquid when it begins to form a crystal lattice & becomes a solid

Heating Curve

Phase Diagram

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