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Particles are packed close together
Particles arranged in organized fashion
Particles have a fixed position
Definite shape and volume
Strong intermolecular attractions

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Particles are farther apart than solids
Particles are free to slide past one another, able to flow
Weak intermolecular attractions
Definite volume and no definite shape

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Particles are very far apart from each other
Particles travel in a random manner
Particles collide with each other and the walls of the container ( pressure )
No intermolecular attractions
No definite shape or volume
Low density
Compressible
Able to flow

 https://www.youtube.com/watch?v=AuOOA
T294vQ

Endothermic (absorbs heat)
Solid → liquid
(melting)
Liquid → gas
(vaporization)
Solid → gas
(sublimation)

Exothermic (releases heat)
Liquid → solid
(freezing)
Gas → liquid
(condensation)
Gas → solid
(deposition)

Refer to your reference packet to answer the following questions. What phase (solid, liquid, or gas) would the following substances exist in at room temperature? (25° C)
1. Methane
Gas
2. Ethanol
Liquid
3. Hydrogen Sulfide
Gas
4. Glucose
Solid
5. Water
Liquid
6. Hexane
Liquid
7. Hydrogen Chloride
Gas


Show the state of a substance at a given temperature and pressure
Liquid
Solid
Critical
Point
Triple
Point
Gas

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
Heat – form of energy
Units – calories (cal) and joules (J)


Temperature – measure of the average kinetic energy
Units ° C, ° F, K



Amount of heat needed to raise the temperature of 1 g of a substance by 1 ° C
Units: J/g ° C or Cal/g ° C

1. What metal has the highest heat capacity?
2. What metal has the lowest heat capacity?
3. If the same amount of heat was added to these metals, which metal would have the largest temperature change?

q = mC p
∆T q = heat C p
= specific heat capacity
∆T = change in temperature

1. How much heat is needed to change 135g of Cu from 20°C to 80°C?

2. Calculate specific heat of a metal if 4501 J of heat were needed to change the temperature of
88.0 g of the metal by 50.0°C. Identify the metal.

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
Heat required to melt 1 g of a substance
Units: J/g
H f
(H
2
O) = 334 J/g
Heat q = mH f mass Heat of fusion

1. How much heat is needed to melt 25.0g of ice?

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
Heat required to vaporize 1 g of a substance
Units: J/g
H v
(H
2
O) = 2260 J/g
Heat q = mH v mass Heat of fusion

1. How much heat is needed to vaporize 250 g of water?

1. Calculate the heat needed to convert 100 g of ice at -20°C to steam at 150°C.

Vapor Pressure:

The pressure of a vapor above a liquid.
Boiling Point:

The temperature at which the vapor pressure of a liquid equals the external pressure.
Boiling:


Occurs throughout liquid
Requires heat
Evaporation:


Only surface molecules
Does not require heat

Rate of evaporation will increase as temperature increases.

Evaporation Boiling
Vapor pressure < Atmospheric pressure
Bubbles cannot form
Vapor pressure = Atmospheric pressure
Bubbles can form and rise


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
Forces between molecules.
Determines whether a substance is a solid, liquid or a gas at room conditions.
Gasno intermolecular forces
Solidstrong intermolecular forces


Three types of forces: dispersion, dipole, H-bonding (hydrogen bonding)





 exist in all molecular substances and atoms.
weakest of intermolecular forces as molecular mass increases, dispersion forces are stronger
Nonpolar molecules have only dispersion forces
Ex: CH
4
, Ar




Found in polar molecules
Stronger than dispersion forces
Ex: PCl
3
, H
2
O

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Hydrogen bonded to a small , highly electronegative atom
*N, O, F*
Stronger than dispersion or dipole-dipole
Ex: H
2
O, NH
3

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N
2 dispersion
CCl
4 dispersion
AsBr
3 dipole
SiCl
4 dispersion
NH
3 hydrogen bonding

 melting/boiling points _______________
 heat of vaporization ________________
 surface tension ____________________
 vapor pressure ____________________




Which of the following compounds has the highest boiling point?
H
2
O, CH
4
, He ____________
Highest vapor pressure? ______________



Which of the following compounds has the highest boiling point?
Kr, Ne, Ar ______________
Highest vapor pressure? ______________

Explain why the boiling point of HF is higher than HCl even though the mass of HF is lower.

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
Liquids that readily evaporate at room temperature are volatile .
Boiling point- temperature at which the vapor pressure of a liquid equals the external pressure
Normal boiling point- external pressure is 1 atm (sea level)

Water: 100°C
Melting point and freezing point are the same temperature

Water: 0°C
Boiling point is affected by altitude
At sea level, air pressure is about 1atm

Atm=atmosphere=unit of pressure


In Denver, at an altitude of about 1600m
(1mile), air pressure= about 0.83atm


The boiling point of H
2
O ≈ 95°C
It takes food longer to cook in Denver

1. What is the vapor pressure of A at 20°C? _____ B? _____ C? ______
2. What is the normal boiling point of A? _____ B? ______ C? _______

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Molecular Solids
Solids formed from covalent compounds.
Forces between molecules are weak.
fairly soft fairly low melting points poor conductors
Ex: sugar, dry ice (CO
2
), ice, S
8
, P
4

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Ionic Solids high melting point conductive in molten form or in solution
Ex: NaCl, CaBr
2

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Metallic Solids good conductors of heat and electricity ductile, malleable luster mobile e (delocalized e )
Ex: Cu, Ag

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Network Covalent Solids
Attached by very strong covalent bonds.
very hard very high melting points poor conductors
Ex: diamond, quartz (SiO
2
)


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
Amorphous solids
No regular pattern
Ex: glass, wax


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
Ice is less dense than liquid water
Ice is a 3 dimensional crystal with holes and empty spaces within the crystal
As H
2 crystal
O freezes, it must expand to form an open
Disadvantages: potholes, pipes bursting
Advantage: when lakes and ponds freeze the ice floats to the top and insulates the water underneath the ice