Chapter 13 chemistry

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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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