Chapter 13 - States of Matter Section 13.1 The nature of Gases
Kinetic energy and gases Kinetic energy: the energy an object has because of its motion Kinetic theory: all matter is made if particles in constant motion There are 3 assumptions: Particles in a gas are small hard spheres with insignificant volume No attractive forces between particles Motions of gas particles are rapid, constant and random In straight lines until collisions, when they rebound. As such they fill any container All collisions between particles are perfectly elastic Energy is transferred from one particle to another without loss. Total kinetic energy remains constant
Gas pressure Gas pressure: pressure that results from force exerted by a gas on an object The amount of force on a given area] This force comes from the collision of gas particles - moving bodies exert a force when they collide with other bodies Vacuum: an empty space with no particles and no pressure No particles - no pressure: i.e. space Atmospheric pressure: due to the collisions of atoms and molecules in air with objects Changes with altitude
How to measure atmospheric pressure A barometer is used to measure atmospheric pressure Particles in the air collide with mercury in the dish, forcing mercury level in the tube to rise Air pressure depends on weather and altitude
Units? The SI unit of pressure is the pascal 1 Pa is very small Normal atmospheric pressure is about 100,000 Pa or 100 KPa Older units are still widespread - millimeters of mercury (mm Hg) or atmosphere One atmosphere is the pressure needed to support 760 mm of mercury in a mercury barometer at 25 C. 1 atm = 760 mm Hg = 101.3 KPa
Kinetic energy and temperature As a substance is heated, particles absorb energy Some is stored in the particles Remaining energy speeds up the particles - increasing kinetic energy and therefore temperature Kinetic energy of particles at a given temperature have a wide range of values Scientists refer to average kinetic energies
Kinetic energy and temperature At any given temperature, the particles of all substances have the same average kinetic energy The physical state does not matter Example - ions in table salt, molecules in liquid water and atoms in gaseous helium all have the same average kinetic temperature at room temperature The average kinetic temperature is directly proportional to the substances temperatures in kelvins Increase in kinetic energy causes temperature to rise As a substance cools, particles tend to move more slowly, The kelvin temperature is directly proportional to the average kinetic energy of the particles of a substance
Section 13.2 The nature of liquids
Particle movements in liquids Particles in both gases and liquids have kinetic energy - allowing particles to flow past one another Substances that contain particles that can flow are called fluids Both gases and liquids - this allows them to take the shape of their containers In liquids - particles are attracted to one another, whereas in gases they are not This is why liquids have a definite volume Properties of liquids are determined by the tendency of particles to flow, and the attraction that keeps the particles close together
Evaporation and Vaporization Vaporization - Conversion of a liquid to a gas Evaporation - When this conversion only occurs at the surface of the liquid, rather than throughout Molecules in a liquid are attracted to one another - particles need a minimum kinetic energy to escape the liquid and turn into a gas
Relationship between evaporation and temperature Evaporation rate increases with temperature Heating a liquid increases the average kinetic energy of particles As evaporation takes place, the liquid s temperature decreases - evaporation is a cooling process This is why sweating is a cooling process
Vapor pressure This is the measure of the force exerted by a gas above a liquid in a sealed container Over time some of the particles may convert from a gas back to a liquid - condensation Eventually the number of particles condensing will equal the number of particles vaporizing - the system reaches an equilibrium An increase in temperature increases the vapor pressure of a system - due to the increased kinetic energy
Measuring vapor pressure Vapor pressure is measure using a manometer The mercury level in the manometer directly relates to vapor pressure
Boiling point When a liquid is heated until particles throughout the liquid have energy to vaporize it has begun to boil The boiling point is when the vapor pressure of the liquid is just equal to the external pressure of the liquid
Relationship between boiling point and pressure Boiling points decrease at higher altitudes In Denver (1600m) water boils at only 95 C The boiling point depends on the external pressure
Boiling point and energy If you continue to heat a liquid, it s temperature will never go beyond the boiling point - it will only boil faster The vapor formed is at the same temperature as the liquid, but with more potential energy (the particles had enough kinetic energy to escape) A steam burn has the potential to cause more damage than a liquid burn Normaling boiling points refer to a standard pressure of 101.3 KPa - atmospheric pressure at sea level
The nature of solids Chapter 13.3
Solids - general characteristics Particles are not free to move Solids are dense and not easy to compress Solids vibrate around fixed points
Melting point vs freezing point The melting point is the temperature that a solid changes into a liquid The freezing point is the temperature that a liquid changes into a solid These are the same temperatures! Ionic solids tend to have higher melting points than molecular solids
Crystal structures and unit cells Most solid structures are crystalline Particles are arranged in an orderly, repeating 3-D pattern The shape depends on the crystals within it Unit cell - the smallest group of particles in a crystal that retain the geometric shape of the crystal
Different crystal systems
Allotropes Two or more different molecular forms of the same element Often highly different properties
Amorphous solids and glasses Not all solids form crystals A solid with randomly arranged atoms is called an amorphous solid Examples - rubber, plastic, asphalt Glasses are a special kind of amorphous solid - the substances cooled without crystallizing The internal structure is intermediate between a crystalline solid and a liquid Glasses do not melt, but soften when heated When glass shatters, fragments have irregular angles and jagged shapes
Changes of State Section 13.4
Sublimation Iodine undergoes sublimation - change of solid to vapor without passing through liquid state Freeze drying process utilises sublimation
Phase diagrams Matter changes between phases due to temperature and pressure Phase diagrams can be used to show the relationship between solid, liquid and gaseous states in a sealed container Lines represent equilibrium between two phases Triple point describes the conditions under which all three phases can exist