States of Matter Water, Solid, Liquid, Gas, and changes in states of matter

Water Found in all three states naturally Can be found on 75 % of Earth’s surface 70 – 90 % of the mass of living things Where many life reactions take place (aqueous)

Water structure H 2 O Has polar- covalent bonds – both are non-metals and the electrons are shared unequally Remember partial charges? Has a 105 degree angle Has two lone pairs Has a tetrahedral molecular geometry Has an sp 3 hybridization- where it bonds

Water’s physical properties and numbers At room temp it is : colorless, odorless, transparent, and tasteless Any disparity is caused by impurities or things dissolved in the water Water freezes and ice begins melting at 0 degrees Celsius Water boils at 100 degrees Celsius Water has a density of g/cm 3 Ice has a density of g/cm 3 Ice has a molar enthalpy of fusion of kJ/mol Water has a molar enthalpy of vaporization of kJ/mol

Hydrogen bonds in water The partial charge on each end allows for strong hydrogen bonds This allows cohesion and adhesion Cohesion- water sticks to itself- enables some bugs to walk on water Adhesion- water sticks to other things- why water forms beads Makes hexagonal patterns as ice and expands up to 9 %- why it floats (otherwise, it would sink and never melt) These bonds and structure give water very high melting point, boiling point, and molar enthalpies compared to similar structures

Solid properties London dispersion forces, dipole-dipole attraction, and hydrogen bonding are all stronger in solids Keeps particles in a fixed position, close together, and with little energy Has a definite shape and volume Has a definite melting point- physical change of solid to liquid High density and incompressible Low rate of diffusion- will not spread easily

Solids- two types Crystalline solids and amorphous Crystalline- particles are arranged in an orderly, geometric, repeating pattern Also called a lattice structure Amorphous- particles are arranged randomly At melting point can be considered supercooled liquid- appear solid, but have liquid properties

Crystals- four types 1. ionic- positive and negative ions in a pattern – hard, brittle, high melting points, and good insulators 2. covalent network- two negative ions- lots of oxides and some transition metals- hard, brittle, high melting points, nonconductors or semiconductors 3. metallic- metal cations and delocalized valence electrons- the electrons belong to the metal but can move around – high electric conductivity 4. covalent molecular- held together by intermolecular forces- so very weak – low melting points, easily vaporized, soft, good insulators

Amorphous solid Means without shape No regular pattern or crystallization Glass and plastic Used in copiers, printers, flat panel displays of computers

liquids Definite volume No definite shape- takes the shape of the container Particles are in constant motion, but fairly close together Lower mobility and intermolecular forces than solids Also called a fluid because of mobility (also refers to a gas) Usu. Flows downhill Except helium, will flow uphill as a liquid In between the solid and gas on properties- the medium level

Liquid properties Relatively high density Relatively incompressibility Able to diffuse Evaporation- physical change to gas without heat Vaporization- physical change to gas from a liquid or solid Boiling- the change of a liquid to bubbles of vapor throughout the liquid Freezing-physical change to a solid

liquids Surface tension- force that pulls all the parts of a liquid’s surface together, keeps the surface area as small as possible creating its properties Capillary action- the attraction of the surface of a liquid to the surface of a solid- pulls the liquid uphill- seen in small tests tubes and plant vascular systems and creates the meniscus you see

Gas properties Expands easily Very fluid- particles always moving Low density Very compressible Diffusion- multiple gases spread out and mix without being stirred Effusion- gases spread out and move through small openings Rates of effusion are directly proportional to the velocities of their particles (the heavier gas molecules move slower)

Gas-two types Real gas- gas that does not behave according to our assumptions Ideal gas- a hypothetical gas the perfectly fits all the assumptions of the kinetic-molecular theory Can be seen at ideal conditions (just the right temp and pressure) in noble gases

Kinetic-molecular theory Based on the idea that particles of matter are always in motion Used to explain the properties and intermolecular forces of solids, liquids, and gases Has five assumptions

Kinetic- molecular five assumptions 1. gases consist of large numbers of tiny particles that are far apart relative to size- explains low density and compressibility 2. collisions between gas particles and between particles and the container was are elastic collisions – no net loss of total kinetic energy as long as the temperature is the same- explains equal but opposite reaction and enforces energy cannot be created or destroyed

Kinetic-molecular five assumptions 3. gas particles are in continuous, rapid, random motion and possess kinetic energy, which is the energy of motion- at the right temperature, gas molecules have so much energy they do not combine with each other 4. there are no forces of attraction between gas particles- when they collide, they do not stick, but bounce apart 5. the temperature of a gas depends on the average kinetic energy of the particles of the gas

Gas equations KE =1/2 mv 2 KE is kinetic energy, m is mass, v is velocity (little v) Speed will increase with temperature increase and decrease with temperature decrease PV= nRT P is pressure, v is volume (big V), n is number of moles, R is a constant number of atm*L/mole*K, T is temperature in Kelvin STP is standard temperature and pressure Pressure is 1 atm (atmosphere) Temperature is K or 0 degrees Celsius Volume is 22.4 liters

example How much energy is possessed by a helium atom moving at 30 m/s ? KE = m v 2 and helium weighs 4 g KE = ½ * 4 (30 2 ) = 1800 m/s

Phase changes

vocab Equilibrium- a dynamic condition in which two opposing changes occur at equal rates in a closed system, no net change Volatile liquids- liquids that evaporate easily, have weak forces of attraction Enthalpy- a measure of energy Molar enthalpy of vaporization- the amount of energy as heat that is needed to vaporize one mole of liquid at its boiling point at a constant pressure, ΔH v Molar enthalpy of fusion- the amount of energy as heat required to melt one mole of solid at its melting point, ΔH f

Molar enthalpy of fusion Energy absorbed or given off as heat when a substance melts or freezes at the melting point of the substance Molar enthalpy of fusion = ΔH f = energy absorbed/ moles of substance

example The molar enthalpy of vaporization for water is kJ/mol. Express this enthalpy of vaporization in joules per gram kJ * 1000 J/kJ = J/mol / 18 g/mol = J/g

Example kJ of energy is required to melt mol of ethylene glycol (C 2 H 6 O 2 ) at its melting point. Calculate the molar enthalpy of fusion, ΔH f, of ethylene glycol and the energy absorbed kJ/ = 11.2 kJ/mol

Example 3 Determine the quantity of energy that will be needed to melt 2.50 x 10 5 kg of iron at its melting point, 1536 degrees Celsius. The ΔH f of iron is kJ/mol kg * 1000 = g/ 56 g/mol Fe = /mol / mol x kJ/ mol = kJ

Phase diagram A graph of P vs T that shows the conditions under which the phases of a substance exist Triple point- the P and T when solid, liquid, and vapor coexist at equilibrium Critical point- the critical P and T, when something will change Critical temperature- the T when it stopes existing as a liquid no matter what the pressure is Critical pressure- the lowest P when it has to be a liquid

Vapor pressure diagram Changes more readily with temperature

example Using the previous graph, estimate the boiling point of ethanol at an applied (external) pressure of 400 torr. (torr is another unit of pressure) roughly degrees Celsius