Chapter 13- States of Matter Flashcards
Kinetic Molecular Theory
3 assumptions
-Describes the behavior of matter in terms of motion and helps explain the properties of gases
Determining Characteristics…
1. Particle Size- gases have small volumes compared to volume of empty space+do not experience attractive or repulsive forces
2. Particle Motion- gases move in constant random motion and their collisions are elastic so they don’t lose energy+kinetic energy is transferred+total kinetic energy does not change
3. Partial Energy- average kinetic energy of a gas is proportional to temp+ as temp increases, velocity and energy increases
Measuring Pressure
- force per unit area
- P=f/a
- 1.00atm=101.325kPa=760mmHg
- pressure is exerted when particles collide with their container’s walls
- more particles colliding= more pressure
- measured with a barometer
- higher elevation= less particles because less gravitational force
Intermolecular Forces
(LDF, DDF, H-Bonding)
- intermolecular forces determine a substance’s state at a given temperature
- dispersion forces are weak forces occurring in all molecules
- dipole-dipole forces happen between opposite charges of polar molecules
- hydrogen bonds are the strongest
Vapor Pressure
-the force of colliding molecules
Viscosity
- measure of the resistance of liquid to flow
- attractive forces slow their movement of particles past one another
AFFECTED BY…
1. attractive forces (>intermolecular forces= >viscosity)
2. particle size+shape (>molecules have >viscosity+slower) (>temp= >kinetic energy)
3. surface tension (extra attraction between molecules at surface of a liquid) (>attraction= >surface tension)
Surface Tension
- extra attraction between molecules at the surface of a liquid
- > attractions between particles= >surface tension
Cohesion+Adhesion
Cohesion- attraction between a molecules and itself
- “like” particles
Adhesion- attraction between a molecule for the molecules in its container
- different particles
Capillary Action
- forces of cohesion and adhesion working together
- example= water in a glass
Crystalline Solids vs Amorphous Solids
Crystalline Solids
4 categories:
1. molecular solids- held together by intermolecular forces+not solid at room temp.+ poor conductors of heat and electricity
2. Covalent Network Solids- composed of atoms that can form multiple covalent bonds (have 4 valence electrons)
3. Ionic Solids- ions are surrounded by ions with opposite charges, high melting points, hard, strong but brittles
4. Metallic Solids- positive metal ions surrounded by a sea of mobile electrons+good conductors of heat and electricity
Amorphous Solids
- “amorphous”=without shape
- solid w/ particles not arranged in geometric pattern
- does not contain crystals
Phase Changes
Main Idea- matter changes when energy is added or removed
- most substances exist in 3 states depending on temp and pressure
- heat= transfer of energy from an object at a high temp to an object of low temp
- boiling point is when molecules have enough energy to vaporize
Phase Changes That Require Energy: (endothermic)
- melting- solid->liquid
- vaporization- liquid->gas/vapor
(evaporation is vaporization at surface of a liquid)
- sublimation- solid->gas
Phase Changes That Release Energy: (exothermic) - freezing- opposite of melting -condensation- opposite of vaporization - deposition opposite of sublimation
Heating Curves
- show how temp changes as a substance is heated up
- describe how the changes in heat take place during phase changes
- plateaus=phase changes because all the molecules are at the same temp
Phase Diagrams
- way to represent phases of substances
- temp and pressure control phase of a substance (have opposite effects so temp increase=vaporization and pressure increase=condensation)
- graphs of temp vs pressure show the phase of a substance under different conditions
