I. States of Matter –Kinetic Molecular Theory –States of Matter

A. Kinetic Molecular Theory KMT –Particles of matter are always in motion. –The kinetic energy (speed) of these particles increases as temperature increases. –Kelvin Temperature scale represents the relationship between temperature and average kinetic energy. K = °C °C = _________ K 23 °C = _________ K 200 K = _________ °C

Evaporation the conversion of a liquid to a vapor below its boiling point What happens… –Molecules at the surface of the liquid go into the vapor state Boiling the conversion of a liquid to a vapor at its boiling point What happens… –Molecules of water vapor form at the bottom and rise to the surface

Boiling Point Definition: the temperature at which the vapor pressure of a liquid is just equal to the external pressure –Bubbles = bubbles of vapor forming throughout the liquid –At lower atmospheric pressures, the boiling point decreases –The temperature of a boiling liquid never rises above its boiling point

Solid Structures –Most solids are crystaline – the particles are arranged in an orderly, repeating, three dimensional pattern.

Amorphous solids lack internal structure – atoms are randomly arranged –Glasses are amorphous solids that are sometimes called supercooled liquids Allotropes are two or more different molecular forms of the same element in the same physical state (ie: diamond or graphite are allotropes of carbon)

Phase Changes Melting/Freezing Vaporization/Condensation Sublimation: Change of a substance from a solid to a gas or vapor w/o passing through the liquid state –Examples: dry ice, iodine Deposition: Change of a substance from gas to solid w/o passing through the liquid state

Phase diagrams Triple point: temperature and pressure at which all three phases are in equilibrium. Critical Point: temp. and pressure past which the liquid and gas phases cannot be distinguished between

Heating Curves Melting - PE  Solid - KE  Liquid - KE  Boiling - PE  Gas - KE 

Copyright 1999, PRENTICE HALL Chapter 1110 Heating Curves Heating Curves

Heating Curves Phase Change –change in Potential Energy (molecular arrangement) –temp remains constant until the phase change is complete Molar Heat of Fusion (  H fus ) –energy required to melt 1 mole of a substance at its melting point –Melting is an endothermic process –+ ΔH

Heating Curves Molar Heat of Solidification –ΔH solid –Energy released when 1 mole of a substance changes from liquid to solid –Exothermic process; -ΔH (heat released) –Heat lost is equal to heat gained during melting

Heating Curves Molar Heat of Vaporization (  H vap ) –energy required to boil 1 mole of a substance at its boiling point –Endothermic; +ΔH

Heating Curves Molar Heat of Condensation –Heat released when one mole of a substance changes from gas to liquid –Exothermic; -ΔH –Heat released is equal to the heat gained during boiling –Ex. Steam burns

Practice Problems How much heat (kJ) is needed to melt 17.0 g of Na? (∆H fus = 2.60 kJ/mol)

1.92 kJ

Given that the molar heat of vaporization of oxygen is 6.82 kJ/mol, how much energy (kJ) would be needed to vaporize g of liquid oxygen?

21.31 kJ

How much heat would be released when 85.0 g of oxygen gas condenses? (∆H vap oxygen is 6.82 kJ/mol, so then what is ∆H condensation )

-18.1 kJ

Example (Heat of combustion) The standard heat of combustion (∆H° rxn ) for glucose (C 6 H 12 O 6 ) is kJ/mol. If you eat and burn 70.g of glucose in one day, how much energy are you getting from the glucose? –Step one: convert g of glucose to moles 70. g glucose x 1 mol = 0.28 mol glucose g –Step two: Use (∆H° rxn ) to find amount of kJ released 0.28 mol glucose x kJ = -790 kJ released 1 1 mol –Step three: if glucose released -790 kJ, then YOU GAINED +790 kJ of energy