States of Matter

Kinetic Theory The word “Kinetic” refers to motion The energy an object has because of it’s motion is called kinetic energy The kinetic theory states that the tiny particles in all forms of matter are in constant motion!

Kinetic Theory & Gases Assumption #1: Three basic assumptions of the kinetic theory as it applies to gases: Assumption #1: Gas is composed of particles- usually molecules or atoms Behave as small, hard spheres Each particle has insignificant volume; relatively far apart from each other (when compared to solids or liquids) No attraction or repulsion between particles (particles are more free to move around more than in a liquid)

Kinetic Theory & Gases Assumption #2: Particles in a gas move rapidly in constant random motion Move in straight paths, changing direction only when colliding with one another or other objects Average speed of O2 in air at 20 oC is an amazing 1700 km/h! (1056 mph)

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Kinetic Theory & Gases Assumption #3: When gas particles collide, no energy is lost. What this means: When large objects (like cars) collide some of the kinetic energy is changed into other forms of energy (such as heat from friction) Before After ECar1 = 10 ECar2 = 10 ECar1 + ECar2 = 15 Ecar1 + ECar2 = 20 + thermal energy (heat) = 5

Kinetic Theory & Gases When particles collide kinetic energy is transferred from one particle to another- the total kinetic energy of the system remains constant When gas particles collide, no energy is lost…this is a perfectly elastic collision After Before Eatom = 10 Eatom = 10 Eatom + Eatom = 20 Eatom + Eatom = 20

Studying the Behavior of Gases Focus on three things: Volume: how much space a gas takes up Pressure Temperature

Gas Pressure Gas Pressure is defined as the force exerted by a gas per unit surface area of an object

Gas Pressure Less collisions Gas exerts pressure because moving particles exert a force when they collide The greater that force, the greater the pressure The greater the number of collisions, the greater the pressure No particles present? Then there cannot be any collisions, and thus no pressure – called a vacuum More collisions Vacuum

Measuring Gas Pressure The SI unit of pressure is the pascal (Pa) Older units of pressure: millimeters of mercury (mm Hg), and atmospheres (atm) Atmospheric pressure: pressure exerted on objects due to the gas particles in the atmosphere. Results from the collisions of air molecules with objects Decreases as you climb a mountain because the air “thins” as elevation increases air density decreases

Barometer measuring device for atmospheric pressure An Early Barometer measuring device for atmospheric pressure which is dependent upon weather & altitude Mercury Barometer – a straight glass tube filled with Hg, and closed at one end; placed in a dish of Hg, with the open end below the surface

Standard Pressure 1 atm = 760 mm Hg = 101.3 kPa An Early Barometer The normal pressure due to the atmosphere at sea level (25oC) can support a column of mercury that is 760 mm high. We call this 1 atmosphere of pressure. 1 atm = 760 mm Hg = 101.3 kPa

Standard Temperature and Pressure Especially for gases, it is important to relate measured values to a standard. This gives us a frame of reference. Standard Temperature and Pressure, or STP --Standard conditions are defined as a temperature of 0 oC and a pressure of 101.3 kPa, or 1 atm Normal boiling point- is defined as the boiling point of a liquid at a pressure of 101.3 kPa (or standard pressure) Normal freezing point, Normal melting point, etc.

Temperature so, temperature increases What happens when you add heat (thermal energy) to a substance? First, particles absorb energy – NO temperature change Some of the energy is stored within the particles as potential energy, and does not raise the temperature Second, temperature increases When the particles cannot store any more energy, additional thermal energy will cause the particles to speed up Average kinetic energy increases so, temperature increases

Water temp. increases 24oC Olive oil temp. increases 50oC How does this happen? Heat and Temperature are not the same thing: Heat is thermal energy, temperature is a measure of the average kinetic energy in molecules Why? The particles in water can store more potential energy Less thermal energy (heat) left over to raise temperature Remember – Water has high specific heat! These are liquids, but gases work the same way. 10g olive oil 10g pure water + 1000 J heat + 1000 J heat Water temp. increases 24oC Olive oil temp. increases 50oC

Average KE & Kelvin Scale The Kelvin temperature scale reflects a direct relationship between temperature and average kinetic energy Particles of a gas at 200 K have twice the average kinetic energy as particles of a gas at 100 K What happens here? When Temperature Kelvin = 0, does kinetic energy reach 0? Doesn’t 0 kinetic energy mean 0 motion? ABSOLUTE ZERO Temp. Kelvin Average Kinetic Energy

The Nature of Liquids Liquid particles are also in motion. Gases and liquids can both FLOW However, liquid particles are attracted to each other, whereas gases are not Intermolecular attractions reduce the amount of space between particles of a liquid. So, liquids are denser than gases (except H2O) Increasing pressure on liquid has hardly any effect on it’s volume

KE of Liquid Particles most of the particles do not have enough energy to escape into the gaseous state; they would have to overcome their intermolecular attractions with other particles Remember: water has hydrogen bonds!

Phase Change (Liquid -> Gas) Vaporization - The conversion of a liquid to a gas or vapor. Vaporization can occur in a liquid that is NOT boiling Vaporization can occur in a liquid that IS boiling

1. Evaporation Evaporation - vaporization occurs in a liquid that is not boiling. Evaporation occurs at the surface of the liquid Some of the particles have more kinetic energy (KE) than others. The particles with enough KE to overcome intermolecular forces break away and enter the gas or vapor state

Effect of Liquid Temperature on Evaporation When you raise a liquid’s temperature by adding heat, the average kinetic energy of the particles is increased which means more liquid particles are likely to have enough kinetic energy to overcome attractive forces between particles (like hydrogen bonds) to become gas particles. So, liquids with higher temperatures will evaporate faster. Particles left behind in liquid are the ones with lower average kinetic energies; they didn’t have enough energy to escape. evaporative cooling - Evaporation is often called a cooling process. On a hot day we sweat to keep cool.

What if the liquid is in a closed container? When some particles do vaporize, these collide with the walls of the container producing vapor pressure Some of the particles will return to the liquid, or condense After a while, the number of particles evaporating will equal the number condensing- the space above the liquid is now saturated with vapor A dynamic equilibrium exists Rate of evaporation = rate of condensation Increasing the temperature increases the KE of the particles  increases the vapor pressure

2. Vaporization of a Boiling Liquid The rate of evaporation from an open container increases as heat is added The heating allows larger numbers of particles at the liquid’s surface to overcome the attractive forces Heating allows the average kinetic energy of all particles to increase

Boiling a Liquid The temperature at which a liquid will boil is called its boiling point (bp). Particles beneath the surface have enough energy to overcome their attractive forces. Gas bubbles form below the surface. At boiling, the vapor pressure of the liquid is just equal to the external pressure on the liquid (atmospheric pressure in an open container) Supplying more heat allows particles to acquire enough KE to escape- the temperature does not go above the boiling point, the liquid only boils at a faster rate

Effect of Pressure on Boiling Point Montrose: Normal bp of water = 100 oC Denver: Normal bp of water = 95 oC, Denver is 1600 m above sea level and average atmospheric pressure is about 85.3 kPa How pressure cookers work: Sealed container allows pressure to increase beyond atmospheric pressure Higher pressure  higher boiling point The temperature rises above 100 oC and food cooks faster.

- Page 394 Normal Boiling Point @ 101.3 kPa = 100 oC Not Boiling Boiling, but @ 34 kPa = 70 oC

Questions:

The Nature of Solids Most solids have particles packed against one another in a highly organized pattern Tend to be dense and incompressible Do not flow, nor take the shape of their container Are still able to move, vibrating about fixed points, rather than moving from place to place (unless they reach absolute zero) Images from www.shorstmeyer.com

Heating Solids When a solid is heated, kinetic energy of the particles increases. More kinetic energy means the particles vibrate more rapidly and more vigorously As a result, solids tend to expand when heated

Phase Change (Solid-> Liquid) If enough heat is added to a solid and the kinetic energy becomes high enough – a solid will melt. The melting point (mp) is the temperature at which a solid turns to a liquid At the melting point, the particle vibrations are strong enough to overcome the interactions holding them in a fixed position Melting can be reversed by cooling the liquid so it freezes Solid liquid

Melting Points of Solids Ionic Compounds and Network Solids: Generally, have high melting points, due to the relatively strong forces holding them together Sodium chloride (an ionic compound) has a melting point = 801 oC Molecular Compounds: Generally, have relatively low melting points because their intermolecular forces are weaker (ex: Vander waals or dipole forces)

Microscopic Structure of Solids Most solid substances are crystalline in structure. In a crystal, the particles (atoms, ions, or molecules) are arranged in a orderly, repeating, three-dimensional pattern called a crystal lattice (All crystals have a regular shape, which reflects their arrangement)

Carbon Solids Carbon is a good example of the significance the crystal structure plays in the compound characteristics. Allotropes are two or more different molecular forms of the same element in the same physical state like carbon. Diamond, Graphite, Buckminsterfullerene are three different substances formed from pure carbon. Differences in how the C atoms bond to each determine the crystal structure and the substance properties.

Phase Change Sublimation- the change of a substance from a solid directly to a vapor, without passing through the liquid state. (solid -> gas) Examples: iodine dry ice (-78 oC); mothballs; solid air fresheners Deposition-the change of a substance from a vapor directly to a solid, without passing through the liquid state (gas -> solid)

Changes of State