Chapter 10 States of Matter 10.1 The Kinetic-Molecular Theory of Matter

Kinetic-Molecular Theory of Gases Particles of matter are ALWAYS in motion; constant, rapid motion. (kinetic energy!) Particles are very small & relatively far apart. Collisions of particles with container walls cause pressure exerted by gas. Volume of individual particles is  zero. Particles exert no attractive or repulsive forces on each other.

Kinetic-Molecular Theory of Gases Gas particles undergo elastic collisions: Collisions in which no energy is lost Average kinetic energy is directly proportional to Kelvin temperature of a gas. Air Hockey Table

Ideal Gas An imaginary gas that perfectly fits all the assumptions of the kinetic-molecular theory A gas with its particles in constant random motion without attraction for each other is called an Ideal Gas. These particles undergo elastic collisions. Nearly all real gases behave as ideal gases EXCEPT at very low temperatures or high pressures.

Likely to behave nearly ideally: Likely not to behave ideally: Real Gases A gas that does not behave completely according to the assumptions of the kinetic-molecular theory. Real gases occupy space and exert attractive forces on one another Likely to behave nearly ideally: Likely not to behave ideally: Gases @ high temp. & low pressure Gases @ low temp. & high pressure Small non-polar gas molecules Large polar gas molecules

Kinetic-Molecular Theory of the Nature of Gases Expansion Gases do not have a definite shape or volume Gases take the shape of their containers Gases evenly distribute themselves within a container Fluidity Gas particles easily flow past one another Low Density A substance in the gaseous state has 1/1000 the density of the same substance in the liquid or solid state Compressibility Gases can be compressed, decreasing the distance between particles, and decreasing the volume occupied by the gas

Kinetic-Molecular Theory of the Nature of Gases Diffusion Spontaneous mixing of particles of two substances caused by their random motion Rate of diffusion is dependent upon: speed of particles diameter of particles attractive forces between the particles

Kinetic-Molecular Theory of the Nature of Gases Effusion Process by which particles under pressure pass through a tiny opening Rate of effusion is dependent upon: speed of particles (small molecules have greater speed than large molecules at the same temperature, so the effuse more rapidly)

Chapter 10 States of Matter 10.2 Liquids

Some Properties of a Liquid Surface Tension: The resistance to an increase in its surface area (polar molecules, liquid metals). A force that tends to pull adjacent parts of a liquid's surface together, thereby decreasing surface area to the smallest possible size.

Some Properties of a Liquid Capillary Action: Spontaneous rising of a liquid in a narrow tube.

Some Properties of a Liquid Viscosity: Resistance to flow (molecules with large intermolecular forces).

Some Properties of Liquids Volatility Liquids that have weak forces of attraction and evaporate easily Nonvolatile Liquids Liquids that have strong forces of attraction and do not evaporate easily

Properties of Fluids Relative High Density 10% less dense than solids (average) Water is an exception 1000x more dense than gases Relative Incompressibility The volume of liquids doesn't change appreciably when pressure is applied Ability to Diffuse Liquids diffuse and mix with other liquids Rate of diffusion increases with temperature

Chapter 10 States of Matter 10.3 Solids

Types of Solids Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS2)].

Types of Solids Amorphous solids aka supercooled liquids: considerable disorder in their structures (glass). Greek for "without shape" Formation of amorphous solids: Rapid cooling of molten materials can prevent the formation of crystals * They do not have definite melting points

Representation of Components in a Crystalline Solid Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

Types of Crystalline Solids Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl).

The smallest portion of a crystal lattice that shows the three-dimensional pattern of the entire lattice Unit Cell

Types of Crystalline Solids Molecular Solid: discrete covalently bonded molecules at each of its lattice points (sucrose, ice).

Packing in Metals Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors.

Closest Packing Holes

Metal Alloys Substitutional Alloy: some metal atoms replaced by others of similar size. brass = Cu/Zn

Metal Alloys (continued) Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. steel = iron + carbon

graphite, diamond, ceramics, glass Network Solids Composed of strong directional covalent bonds that are best viewed as a “giant molecule”. brittle (non-flexible) do not conduct heat or electricity carbon, silicon-based graphite, diamond, ceramics, glass

Sulfur – S8

Phosphorus – P4

Diamond

Graphite

Zirconia

Chapter 10 States of Matter 10.4 Changes of State

Equilibrium Dynamic condition in which two opposing changes occur at equal rates in a closed system A closed system at constant temperature will reach an equilibrium position at which the rates of evaporation and condensation will be the same

Equilibrium Vapor Pressure The pressure of the vapor present at equilibrium. Determined principally by the size of the intermolecular forces in the liquid. Increases significantly with temperature. Volatile liquids have high vapor pressures. Increasing the temperature will move more particles into the vapor phase to compensate for the new energy

Boiling The conversion of a liquid to a vapor within the liquid as well as at its surface. It occurs when the equilibrium vapor pressure of the liquid equals the atmospheric pressure Boiling Point The temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure Water boils at 100 °C at 1 atm pressure Water boils above 100 °C at higher pressures Water boils below 100 °C at lower pressures

LeChatelier’s Principle When a system at equilibrium is placed under stress, the system will undergo a change in such a way as to relieve that stress.

Translation: When you take something away from a system at equilibrium, the system shifts in such a way as to replace what you’ve taken away. When you add something to a system at equilibrium, the system shifts in such a way as to use up what you’ve added.

LeChatelier’s Example #1 A closed container of ice and water at equilibrium. The temperature is raised. Ice + Energy  Water The equilibrium of the system shifts to the _______ to use up the added energy. right

LeChatelier’s Example #2 A closed container of N2O4 and NO2 at equilibrium. NO2 is added to the container. N2O4 + Energy  2 NO2 The equilibrium of the system shifts to the _______ to use up the added NO2. left

LeChatelier’s Example #3 A closed container of water and its vapor at equilibrium. Vapor is removed from the system. water + Energy  vapor The equilibrium of the system shifts to the _______ to produce more vapor. right

Water phase changes constant Temperature remains __________ during a phase change. constant

Phase Diagram Represents phases as a function of temperature and pressure. Critical temperature: temperature above which the vapor can not be liquefied. Critical pressure: pressure required to liquefy AT the critical temperature. Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm).

Water Water

Phase changes by Name

Carbon dioxide Carbon dioxide

Carbon Carbon

Sulfur

Chapter 10 States of Matter 10.5 Water

Water’s Properties

Sea Ice Ice forms on top of the ocean in a thin layer & acts to insulate the warmer waters below from the colder air temperatures. This occurs in the polar regions, the Artic & Antarctic Since ice is less dense than liquid water, it will float on top, instead of sinking which would kill all life below the surface. Sea ice is not the same as an iceberg. Icebergs are pieces of glaciers which are formed by snowfall on land. Sea ice is not salty, as the hydrogen bonds that hold ice together will not form properly if salt remains in the structure.

Temperature & Salinity Sea Water Density There are 2 main factors that affect the density of sea water: Temperature & Salinity As the temperature decreases, density increases. As salinity increases, density increases. Since the least dense layer of a liquid will float above a more dense layer, the warmest & lowest salinity layer will be on top. However, temperature has a greater effect on density than salinity. That means that a higher salinity layer can float on top of a lower salinity layer if it is considerably warmer in temperature.

Molar Heat of Fusion The amount of heat energy required to melt one mole of solid at its melting point. 6.009 kJ per 1 mole Practice! How much energy is absorbed when 16.3 g of ice melts? 16.3 g x 1 mole x 6.009 kJ = 5.44 kJ 18 g 1 mol

Molar Heat of Vaporization The amount of heat energy required to vaporize one mole of a liquid at its boiling point Strong attractive forces between particles result in high molar heat of vaporization 40.79 kJ per 1 mole Practice! Find the mass of liquid water required to absorb 5.23 x 104 kJ of energy upon boiling. 5.23 x 104 kJ x 1 mole x 18 g = 2.31 x 104 g 40.79kJ 1 mol

Specific Heat The amount of energy required to change the temperature of 1 gram of a substance by 1 degree Celsius is known as specific heat capacity. (Ch.16 p.533) Q = s * m * ΔT Where Q = energy (heat) required s = specific heat capacity m = mass of sample, grams ΔT = change in temperature, Celsius

Specific Heat Capacities of Some Common Substances Specific Heat Capacity (J/g°C) Water, Liquid 4.184 Water, Ice 2.03 Water, Steam 2.0 Aluminum, s 0.89 Iron, s 0.45 Mercury, l 0.14 Carbon, s 0.71 Silver, s 0.24 Gold, s 0.13

Practice! A piece of copper alloy with a mass of 85.0 g is heated from 30.°C to 45°C. In the process, it absorbs 523 J of energy as heat. What is the specific heat for this copper alloy? How much energy will the same sample lose if it is cooled from 45°C to 25°C? Answers… 0.41 J/g∙K 7.0 x 102 J

Specific Gravity A function of density. The ratio of density of a material to the density of water at a specified temperature. The density of water is usually 1 g/cm3 When dividing density by density, the units cancel and therefore, specific gravity has no units! If the density of the material is in the units g/cm3 (or g/mL) when dividing by the density of water, the numeric value of the material’s density is the same, but the units cancel out.

Intermolecular Forces (Review from Chapter 6) Forces of attraction between different molecules rather than bonding forces within the same molecule. Dipole-dipole attraction Hydrogen bonds Dispersion forces

Dipole-Dipole Attraction dipole-dipole attraction: molecules with dipoles orient themselves so that “+” and “-” ends of the dipoles are close to each other.

Dipole Forces

Hydrogen Bonding hydrogen bonds: dipole-dipole attraction in which hydrogen is bound to a highly electronegative atom. (F, O, N)

Hydrogen Bonding in Water

Hydrogen Bonding in DNA

London Dispersion Forces Dispersion forces: relatively weak forces caused by instantaneous dipole, in which electron distribution becomes asymmetrical. The ONLY forces of attraction that exist among noble gas atoms and nonpolar molecules. (Ar, C8H18)

London Dispersion Forces

Boiling point as a measure of intermolecular attractive forces