Ch. 13 States of Matter

Section 1: Gases I. Kinetic-Molecular Theory: explains the properties of gases in terms of the energy, size, and motion of the particles A. Particle size 1. Small volume of particles/large volume of empty space 2. Far enough apart that there are no attractive or repulsive forces among gas particles

B. Particle Motion 1. gas particles are in constant, random motion 2. collisions between gas particles are elastic collisions- no kinetic energy is lost, only transferred. The particles have no influence on each other. C. Particle Energy 1. mass is the same for all particles but the velocity is not; KE is not the same

2. Temperature is a measure of the average KE of the particles, so at a given temp., all gases have the same average KE. II. Behavior of Gases A. Low density: in a large amount of space (volume), the mass of the particles is small; so the DENSITY is small (D = m / v) B. Compression & expansion: the large amount of empty space between gas particles allows the particles to be pushed into a smaller volume

C. Diffusion: movement of one material through another due to the random motion of the gas particles D. Effusion: a gas escapes through a tiny opening; the rate of effusion is inversely proportional to the square root of its molar mass (the heavier the gas particles, the more slowly they diffuse) III. Gas Pressure A. Pressure: force per unit area P = F/A 1. Gas particles exert pressure when they collide with the walls of their container (see pg. 388) 1. Gas particles exert pressure when they collide with the walls of their container (see pg. 388)

2. air particles exert pressure in all directions and produce air or atmospheric pressure; higher altitudes have slightly lower air pressure than lower altitudes 2. air particles exert pressure in all directions and produce air or atmospheric pressure; higher altitudes have slightly lower air pressure than lower altitudes B. Measuring Air Pressure 1. Two instruments used to measure air pressure are 1. Two instruments used to measure air pressure are a. Barometer – invented by Torricelli a. Barometer – invented by Torricelli b. Manometer – measures gas pressure in a closed container b. Manometer – measures gas pressure in a closed container ****The two forces that determine the exact height of mercury in a barometer are gravity and air pressure ****The two forces that determine the exact height of mercury in a barometer are gravity and air pressure

2. Changes in air pressure can be caused by air temperature and humidity. 2. Changes in air pressure can be caused by air temperature and humidity. C. Units of pressure 1. kilopascal (kPa) 2. millimeters of mercury (mm Hg) 3. torr 4. pounds per square inch (psi) 5. atmosphere (atm)

F.Y.I. !!! The SI unit of pressure is the pascal (Pa) The SI unit of pressure is the pascal (Pa) At sea level (normal atmospheric pressure), the average air pressure (at 0 o ) is At sea level (normal atmospheric pressure), the average air pressure (at 0 o ) is 1 atm = 760 mm Hg = 760 torr = kPa Let’s look at the “problem-solving LAB” on pg. 390.

D. Dalton’s Law of Partial Pressures 1. each gas in a mixture exerts pressure independently of the other gases present 2. His 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. P total = P 1 + P 2 + P 3 + P 4 + …..P n P total = P 1 + P 2 + P 3 + P 4 + …..P n The pressure contributed by each P n is called the partial pressure.

3. (#4) P H = ? 600 = P H 600 = P H P H = 161 mm Hg P H = 161 mm Hg (#5) P total = ? P total = P total = P total = kPa P total = kPa (#6) P CO2 = ? 30.4 = P CO = P CO2 P CO2 = 10.2 kPa P CO2 = 10.2 kPa

Sec. 2: Forces of Attraction I. Intramolecular Forces (bonding forces): The attractive forces that hold particles together in ionic, covalent, and metallic bonds (“within”) A. Ionic --- cations and anions B. Covalent --- positive nuclei and shared electrons C. Metallic --- metal cations and mobile electrons

II. Intermolecular Forces: the forces of attraction of particles “between” or “among” other particles; EX: water molecules in a drop of water; weaker than intramolecular (bonding) forces. A. Dispersion Forces (nonpolar mixtures) 1. weak forces that result from temporary shifts in the density of electrons in electron clouds (London forces) – see fig weak forces that result from temporary shifts in the density of electrons in electron clouds (London forces) – see fig temporary dipoles are formed (pg. 394) 2. temporary dipoles are formed (pg. 394) 3. weakest type of intermolecular force; can only be the dominant force of attraction when there are no stronger forces of attraction acting on the particles 3. weakest type of intermolecular force; can only be the dominant force of attraction when there are no stronger forces of attraction acting on the particles

B. Dipole-dipole forces 1. attractions between oppositely charged regions of polar molecules 2. neighboring polar molecules orient themselves to that oppositely charged regions line up (see fig. 13-9) 3. are stronger than dispersion forces if the molecules being compared have approximately the same mass

C. Hydrogen Bonds 1. dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a small, highly electronegative atom with at least one lone electron pair 2. hydrogen must be bonded to either a F, O, or N to form 3. the hydrogen bonds between water molecules are stronger than those between ammonia molecules because oxygen has a higher electronegativity than nitrogen (see fig )

Sec. 3: Liquids and Solids I. Liquids: Characteristics A. Density and compression 1. liquids are MUCH denser than gases 1. liquids are MUCH denser than gases 2. when liquids are compressed, the changes in volume is MUCH smaller than for gases because the liquids are already tightly packed together 2. when liquids are compressed, the changes in volume is MUCH smaller than for gases because the liquids are already tightly packed together

B. Fluidity: the ability to flow 1. liquids diffuse more slowly than a gas at the same temperature 1. liquids diffuse more slowly than a gas at the same temperature 2. liquids are less fluid than gases because of the intermolecular forces 2. liquids are less fluid than gases because of the intermolecular forces C. Viscosity 1. a measure of the resistance of a liquid to flow 1. a measure of the resistance of a liquid to flow 2. determined by the type of intermolecular forces involved, the shape of the particles, and the temperature 2. determined by the type of intermolecular forces involved, the shape of the particles, and the temperature

D. Viscosity and temperature 1. viscosity decreases with temperature 1. viscosity decreases with temperature 2. higher temps. make it easier for the molecules to overcome the intermolecular forces that keep the molecules from flowing (motor oil) 2. higher temps. make it easier for the molecules to overcome the intermolecular forces that keep the molecules from flowing (motor oil) E. Surface tension 1. the energy required to increase the surface area of a liquid by a given amount; a measure of the inward pull by particles in the interior 1. the energy required to increase the surface area of a liquid by a given amount; a measure of the inward pull by particles in the interior

2. the stronger the attractions between particles, the greater the surface tension 2. the stronger the attractions between particles, the greater the surface tension ***surfactant: compounds that lower the surface tension of water F. Capillary action 1. cohesion – the force of attraction between identical molecules 1. cohesion – the force of attraction between identical molecules 2. adhesion – the force of attraction between molecules that are different 2. adhesion – the force of attraction between molecules that are different (see fig on pg. 399) (see fig on pg. 399)

II. Solids: Characteristics A. Density of solids 1. the particles in a solid are more closely packed than those in a liquid 1. the particles in a solid are more closely packed than those in a liquid 2. most solids are more dense than most liquids (exception: ice and liquid waters— see fig , pg. 400) 2. most solids are more dense than most liquids (exception: ice and liquid waters— see fig , pg. 400) B. Crystalline solids– solid whose atoms, ions, or molecules are arranged in an orderly, geometric, 3-D structure Types of crystalline solids 1. atomic solids (Gr. 8A elements) 1. atomic solids (Gr. 8A elements) a. made of atoms a. made of atoms b. Soft to very soft; very low melting points; poor conductivity

2. molecular solids (F 2, H 2 O, NH 3 ) a. Made of molecules a. Made of molecules b. Fairly soft; low to mod. High melting points; poor conductivity b. Fairly soft; low to mod. High melting points; poor conductivity 3.covalent network solids (C & Si cpds) a. atoms connected by covalent bonds and usually have multiple covalent bonds a. atoms connected by covalent bonds and usually have multiple covalent bonds b. Very hard; very high melting points; usually poor conductivity b. Very hard; very high melting points; usually poor conductivity

4. ionic solids (NaCl, CaCO 3 ) a. Made of ions a. Made of ions b. Hard; brittle; high melting points; poor conductivity b. Hard; brittle; high melting points; poor conductivity 5. metallic solids a. Atoms surrounded by mobile valence electrons a. Atoms surrounded by mobile valence electrons b. Soft to hard; low to very high melting points; malleable and ductile; excellent conductivity b. Soft to hard; low to very high melting points; malleable and ductile; excellent conductivity

6. amorphous solids (NOT a crystalline solid) a. A solid in which the particles are not arranged in a regular, repeating pattern (“without shape”) a. A solid in which the particles are not arranged in a regular, repeating pattern (“without shape”) b. Forms when a molten material cools too quickly to allow enough time for crystals to form b. Forms when a molten material cools too quickly to allow enough time for crystals to form

Sec. 4: Phase Changes: usually depend on the temperature and pressure I. Phase changes that require energy A. Melting 1. heat is the transfer of energy from an object at a higher temp. to an object at a lower temp. 1. heat is the transfer of energy from an object at a higher temp. to an object at a lower temp. 2. the energy absorbed by the ice is NOT used to raise the temperature of the ice but to disrupt the H-bonds holding the water molecules together and move apart 2. the energy absorbed by the ice is NOT used to raise the temperature of the ice but to disrupt the H-bonds holding the water molecules together and move apart 3. When the ice absorbs enough energy to break the H-bonds, they move apart and enter the liquid phase 3. When the ice absorbs enough energy to break the H-bonds, they move apart and enter the liquid phase

B. Vaporization 1. the process by which a liquid changes to a gas or vapor 1. the process by which a liquid changes to a gas or vapor 2. when vaporization occurs only at the surface of a liquid, it is called evaporation 2. when vaporization occurs only at the surface of a liquid, it is called evaporation 3. Evaporation is a cooling process because it pulls in energy from the surroundings!! 3. Evaporation is a cooling process because it pulls in energy from the surroundings!!

C. Sublimation 1. process by which a solid changes directly to a gas without first becoming a liquid (no puddles!) 1. process by which a solid changes directly to a gas without first becoming a liquid (no puddles!) 2. examples are dry ice, moth balls, and solid air fresheners 2. examples are dry ice, moth balls, and solid air fresheners 3. Ice sublimes in a shorter time period at extremely low pressure -- called freeze-drying. 3. Ice sublimes in a shorter time period at extremely low pressure -- called freeze-drying.

II. Phase changes that release energy A. Condensation 1. when a molecule loses energy (releases it), its velocity is reduced 1. when a molecule loses energy (releases it), its velocity is reduced 2. process by which a gas or vapor becomes a liquid 2. process by which a gas or vapor becomes a liquid 3. examples include dew, fog, and clouds 3. examples include dew, fog, and clouds B. Deposition 1. a substance changes from a gas or vapor to a solid without first becoming a liquid 1. a substance changes from a gas or vapor to a solid without first becoming a liquid

2. the reverse of sublimation 2. the reverse of sublimation 3. examples: snowflakes and frost 3. examples: snowflakes and frost C. Freezing 1. molecules lose KE and their velocity decreases and the molecules are less likely to flow past each other 1. molecules lose KE and their velocity decreases and the molecules are less likely to flow past each other 2. freezing point: the temp. at which a liquid is converted into a crystalline solid 2. freezing point: the temp. at which a liquid is converted into a crystalline solid 3. the melting point and freezing point are the same for a given substance 3. the melting point and freezing point are the same for a given substance

III. Phase Diagrams A. Temperature and pressure have opposite effects on a substance Ex: an increase in temp. will cause a liquid to vaporize, but an increase in pressure will cause the same liquid to condense. B. A phase diagram is a graph of pressure versus temperature that shows in which phase a substance exists under different conditions of temperature and pressure.

C. Triple point -- the point on a phase diagram that represents the temperature and pressure at which three phases of a substance can coexist; all 6 phase changes occur at the triple point D. Critical point -- indicates the critical pressure and critical temp. above which a substance cannot exist as a liquid

Yellow Yellow curve = solid and vapor exist Blue Blue curve = liquid and vapor exist Red Red curve = solid and liquid exist Triple point = solid, liquid, and vapor exist