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Chapter 1- States of Matter
Gases
Kinetic Molecular Theory
· Model to describe properties of gases
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· Kinetic=move
· Theory describes behavior in terms of particles in motion
· Model makes assumptions about size, motion, and energy of gas particles
Particle size
· Gases consist of small particles that are separated from one another by empty space
· Because gas particles are far apart, there are no significant attractive or repulsive forces among them
Particle Motion
· Gas particles are in constant random motion collide with other particles or with the walls of their container
· Have elastic collisions-no kinetic energy is lost
· Kinetic energy can be transferred between colliding particles, but the total kinetic energy of the particles don't change
Particle energy
· factors determine the kinetic energy of a particle
· Mass, velocity KE=1/mv
· In a sample of a dingle gas all particles have the same mass but all particles do not have the same velocity
· Because of this particles do not have the same kinetic energy
· Kinetic energy and temperature are related
· Temperature- a measure of the average kinetic energy of the particles of matter
· At a given temperature, all gases have the same average kinetic energy
Explaining the behavior of Gases
Low Density
· Theory states- great deal of space exists between gas particles
· The rest are fewer chlorine molecules than gold atoms in the same volume Cl=solid Au=gold
Diffusion & Effusion
Theory states-no force of attraction between gas particles
· Gas particles can move easily past each other
· The space into which a gas flows is already occupied by another gas
· This random motion causes gases to mix until they are evenly distributed
· Diffusion- Movement of one material through another
· Move from high concentration to low concentration
· The rate of diffusion depends on the mass of the particles
· For lighter particles to have the same average kinetic energy as heavier particles, they must have a greater velocity (KE=1/mv
· Effusion is process related to diffusion
· Graham's Law of Effusion- States that the rate of effusion for a gas is inversely proportional to the square root of the molar mass (GFM)
· This law applies to rates of diffusion
· Using this law you can setup a proportion to compare the diffusion rates for gases
Formula
Gas pressure
· Pressure-defined as force per unit area
· Gas particles exert pressure when they collide with the walls of their container
Atmospheric pressure
· Because the particles in air move in every direction, they exert pressure in all directions
Measuring Air pressure
· Barometer-instrument used to measure atmospheric pressure
· Height of the Hg in the tube is determined by forces
· Gravity pushing down on Hg; Humidity/temperature causes air pressure to change
· An increase in air pressure causes Hg to rise, a decrease it falls
· Monometer-instrument used to measure gas pressure in a closed container
· Flask in connected to a U-tube containing Hg
· Before the gas is released Hg is at the same height in arms
· After the gas is released the height in the arms are no longer equal
· If gas pressure atmospheric pressure the liquid on the left decreases and the liquid on the left increases Pgas=Patm+Phg
· If gas pressure atmospheric pressure the liquid on left increases and liquid on right decreases Pgas=Patm-PHg
Units of Pressure
Pascal (Pa)-Si unit of pressure 1Pa=1N/M
Other units mmHg, torr, atmosphere (atm), psi, kilopascal (kPal)
1.00atm=760 torr=101. kPa=14.7 psi (at 0 C)
Dalton's Law of Partial Pressure
· Law states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture
· Partial pressure depends on Number of moles of gas, size of the container, temperature of mixture
· It doesn't depend on the identity of the gas Patm= p1 + p + p….
· When gases are made in a lab they are made over water, so the gas is saturated with water
· Pgas=Ptotal-Pwater
· Water pressure will be given to you Practice problems on pg 4-6
· Water pressure will be given to you.
Forces of attraction
· Intramolecular force-Force within atoms, ions, molecules (bonding)
· Ex dispersion, H-bonds
· Intermolecular force-force between atoms, ions, or molecules
· Ex dispersion, H-bonds
Dispersion force
· Weak forces that result from temporary shift of electrons in the electron cloud
· Occurs when non polar molecules are in close contact or collide
· Electrons in the cloud repel each other causing the electrons to have a higher density in one region of each cloud
· Forms temporary dipole; When the dipoles are close together a weak dispersion force exists; Occurs in identical non polar molecules
Dipole-Dipole forces
· Attractions between oppositely charged regions of polar molecules
· Contain permanent dipoles
· Neighboring polar molecules orient themselves so that charged regions line up.
· Forces are usually stronger than dispersion forces
· Pg 4-5 pictures of bonds
Hydrogen bonds
· A dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a small electronegative atom with at least one lone electron pair
· Hydrogen must be bonded to Fluorine, Oxygen, Nitrogen
· H has a positive charge, O has a negative charge so that one molecule is attracted to the O in another molecule
Liquids
Kinetic Molecular Model
· Have a constant motion
· No foxed position so they take the shape of a container
· Force of attraction is greater than in gases so liquids have a fixed volume
Density & Compression
· At 5C and 1atm, liquids are closer than gases because of their greater intermolecular forces holding them together
· Like gases, liquids can be compressed, but only by a small volume because the molecules are relatively close together
Fluidity
· Ability to flow
· Liquids can diffuse through other liquids but at a slower rate than gases because of intermolecular forces
· Liquids are less fluid than gases
Viscosity
· Is the measure of the resistance of a liquid to flow
· The attractive force slows their movement
· Viscosity is determined by the type of intermolecular forces, shape of the particle, and temperature
· Warm-faster flow, cooler- slow flow larger/longer-slow small/short-fast
Surface tension
· Particles in the middle of a liquid can be attracted to particles all around them
· Particles at the surface have no attractive force from above to balance the attraction for below; so there is a greater force pulling down on particles at the surface
· Surface tension- energy required to increase the surface area of a liquid by a given amount
· Generally the stronger the attraction between particles the greater the surface tension
· Water has a high surface tension because it can form multiple hydrogen bonds
· Surfactants- compound that lower the surface tension of water (break hydrogen bonds)
Capillary action
· A narrow container; surface of water forms a meniscus forces cause this
· Cohesion- force of attraction between identical molecules
· Adhesion- force of attraction between different molecules
· The adhesive force between the water and glass is greater than the cohesive force between the water; so water rises along the inner walls of the cylinder
Kinetic molecular theory
· Are in constant motion
· Have a strong attractive force between particles, this limits the motion, causing more order
· Because this solids have a definite shape and volume are much less fluid
Density of solids
· Particles are packed closely together so they are more dense
· Because most solids sink, except water has fewer particles in given volume
Crystalline solids
· Solid where atoms, ions, or molecules are arranged in a orderly, geometric shape
· Unit cell- smallest arrangement of connected points that can be repeated in 1 direction to form a lattice
· Shape of a crystal is determined by the type of unit cell
· 7 different crystals on pg 401
· Crystalline solids can be classified into 5 groups based on the type of particles they contain
· Pg 40 table 1-4
· Amorphous solids- particles are not arranged in a regular repeating pattern
· Forms when the motion material cools too quickly and the crystals don't have time to form
Phase changes
MeltingÞsolidÞliquid
· Solids well absorb heat energy and break the bonds that are holding them together
· Temperature does not change as the solid is melting
· Amount of energy needed depends on the strength of that bond
· Melting point- temperature at which the liquid phase and solid phase of a given substance can coexist
Vaporization liquidÞgas
· Once all ice melts the addition of energy increases the kinetic energy of the liquid and the temperature increases
· Evaporation- when vaporization occurs only at the surface of a liquid
· Liquid molecules must absorb energy, break forces of attraction, and enter gas phase ex sweat
· In a closed container water collects above that liquid and exerts pressure on the surface of the liquid called vapor pressure
· Boiling point- temperature at which vapor pressure of a liquid equals the external or atmospheric pressure
Sublimation solidÞgas
· Ex solid iodine, solid carbon dioxide, mothballs, air freshener
· Absorb energy
Condensation gasÞliquid
· Loses energy, velocity decreases when molecules collide they are more likely to form bonds or increase the force of attraction, and changes to the liquid phase
· Reverse vaporization
Deposition gasÞsolid
· Reverse of sublimation energy is released ex frost & snow
Freezing liquidÞsolid
· Energy is lost, velocity decreases bonds form and hold the molecules in a fixed position
· Reverse of melting
· Freezing point- temperature at which a liquid is converted in to a solid
· Is equal to the melting point
Phase diagrams
· Graph of pressure vs. temperature that shown in which phase a substance exists
· Under different conditions of temperature and pressure
· curves separate the regions from one another solid, liquid gas
· On the curve phases of matter exist
· Triple point-temperature at which 6 phase changes can occur and phases of matter can exist
· Critical point- indicates the critical temperature and pressure above which a substance cannot exist as a liquid
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