Transcript Ch 10 Notes_Teacher

Liquids and Solids
Ch 10.2,10.3 & 10.4
Pg. 353 # 4, 5, 7, 8, 10-14,17, 20-22, 27,
28, 33
Liquids exist in the smallest temperature range, so
liquids are the least common state of matter . . .
 Kinetic Theory Description of the Liquid State

According to the kinetic theory, motion of liquid
particles can be described as . . . Constant.
However, the particles in a liquid are closer
together than a gas would be, but they are not
bound in a fixed position.
Properties of Liquids and the Particles Model
– define each property Properties of Fluids
 Definite Volume – particles have
enough attractive forces to keep them
close together. They cannot expand to
fill a container
 Fluidity – ability to flow (or move
past each other) and take shape of
container
 Relative High Density – close
arrangement of particles (compared to
a gas) making mass/volume ratio
higher.
 Incompressible – much less
compressible (can’t be squeezed
together) than gases b/c particles are
closer together
 Dissolving Ability – see ability to
diffuse
 Ability to Diffuse – can mix with
other liquids due to constant,
random motion of particles (slower
than gases)
 Surface Tension –a force that
pulls adjacent parts of the liquid’s
surface together, thereby decreasing
surface area. Results from the
attractive forces between particles.
 Tendency to Evaporate
and Boil – Vaporization is the
liquid to gas phase change.
 Tendency to Solidify –
Freezing is the liquid to solid phase
change
10.2 Questions
 Why are liquids more dense than gases?

Particles are closer together so more particles are in a given area
 Why are liquids harder to compress than gases?

Particles are closer together so more particles are in a given area.
There is less space between particles so less room for
compression.
 Why do liquids diffuse slower than gases?

Liquid particles are closer together. Attractive forces between
particles slow their movement.
 Can a liquid boil without increasing the
temperature? How?

Yes – lower the atmospheric pressure
10.3 Solids
 “Solid as a rock, “ is the description of solid –
something that is hard, unyielding, with a definite
shape and volume. Many things other than rocks are
solids. In fact, solids are more common than liquids.
This diagram shows the particles of a gas, liquid and solid.
Kinetic-Theory Description of the Solid State
According to the kinetic theory, the motion of solid particles
can be described as….
 Lower kinetic energy, less motion, more packed particles,
and higher intermolecular forces (IMF)
 Properties of Solids and the Particle Model – define each property:
 Definite shape and volume – solids maintain a definite shape without a
container. Volume is constant due to closely packed particles.
 Non-fluid – particles can’t flow because particles are held in relatively
fixed positions.
 Definite melting point – The temperature at which the kinetic energy of
the particles are able to overcome the attractive forces holding them
together in fixed positions (crystalline only)
 High Density – solids are packed more closely than that of a liquid or
gas.
 Incompressible – particles are packed so close together there is
virtually no space between them.

Slow Diffusion – Millions of times slower than liquids due to the
high IMF’s between particles.
Crystalline Solids
 Classification of crystals by arrangement
and shape
 Crystal Lattice (crystal structure) – The total 3-D
arrangement of particles of a crystal.

The arrangement of particles in the crystal can be
represented by a coordinate system called a lattice.
 The smallest portion of the crystal lattice that reveals
the 3-D pattern of the entire lattices is the unit cell.
Binding Forces in Crystals
Simple
Body-centered
(ex. Li, K, Cr)
Types of Crystals
Face-centered (ex.
Cu, Ag, Au)
Hexogonal (like oranges
in
a grocery store); (ex. Zn)
Binding forces in crystals
Binding Force
Ionic crystals
Covalent network
crystals
Metallic crystals
Covalent molecular
crystals
Lattice consists of
Formed When /
Binding Force
(+) and (-) ion as
arranged in regular
patterns
Group 1/2 metals
combine with Group
7/8 nonmetals
single atoms
Atoms covalently bond
to neighbors,
extending through a
large number of
atoms (chains form)
(+) ions of the metal
surrounded by a
sea of valence
electrons
Each e- and the (+)
metallic ions attract
Covalently bonded
molecules held
through IMF
For nonpolar
molecules, London
Forces; For polar
molecules, DipoleDipole.
Amorphous Solids
 Rubber, glass, plastics and synthetic fibers
are called amorphous solids.
“Amorphous,” comes from the Greek
for “without a shape.”
 Unlike crystals, amorphous solids do not
have a regular, natural shape, but instead
take on whatever shape imposed on them.
 Particle arrangement is not uniform; they are
arranged randomly, like particles of a liquid.
 Examples of amorphous solids – rubber, wax,
plastics, synthetic fibers, glass used in fiber
optics (optical fibers transmit telephone
conversations by means of light waves.
Amorphous solids are prepared by cooling molten
materials in a way that prevents crystallization..
 Molecular examples
Crystalline vs. Amorphous
10.4 Changes of State
Possible Changes of State
Change of State
Name
Example
Solid -> Liquid
melting
ice -> water
Liquid -> Solid
freezing
water to ice
Liquid -> Gas
vaporization
Br(l) -> Br(g)
Gas -> Liquid
condensation
water vapor -> water
Solid -> Gas
sublimation
dry ice -> CO2 gas
Gas -> solid
deposition
frost
Equilibrium
 What does equilibrium mean?
 It is a dynamic condition in which two
opposing changes occur at equal rates in a
closed system.
 What is a closed system?
 A container with a lid
 When a liquid changes to a vapor, as in
evaporation, it absorbs heat energy and can be
shown as:

Open system evaporation –

Closed system evaporation –
liquid + heat  vapor
liquid + heat vapor
When a vapor condenses, as in condensation,
it gives off heat energy and can be shown as:

And condensation –
vapor  liquid + heat
 The liquid vapor equilibrium can be rewritten as:
↔ vapor

liquid + heat

“The double yields sign represents a reaction at
equilibrium”
Le Chatelier’s Principle
 What is it? When a system at equilibrium is
disturbed by the application of stress, the
system reacts to minimize the stress.
 Is temperature an example of stress?

Yes.
 What happens when you increase the
temperature of a system? Equ. Shifts away
from heat
 ↓ liquid + increased heat ---> ↑ vapor
Le Chatelier’s Principle
 What happens when you decrease the
temperature of a system? Equ. Shifts toward
heat!

↓ vapor
---> ↑ liquid + decreased heat
 What factor is controlling the decrease and
increase of vapor and liquid?

the temperature (heat)
Equilibrium Vapor Pressure of a Liquid
 What is it?

At equilibrium, the molecules of a vapor exert
a specific pressure on its corresponding
liquid.
When equilibrium vapor pressure of water is
graphed, (draw figure 14 below):
 The strength of attractive forces is
independent of temperature. Higher
temperatures with resultant higher kinetic
energies make these forces less effective.
 Liquid water can exist in equilibrium with
water vapor only up to a temperature of
374.1ºC. Later you will learn that neither
liquid water nor water vapor can exist at
temperatures above 374.1ºC.
Water
Alcohol
Cooking Oil
At
80° C
355 torr
760 torr
10 torr
At
50° C
92 torr
400 torr
4 torr
At
20° C
20 torr
90 torr
1 torr
 What is equilibrium called when liquid
molecules enter into the gaseous state?

Vaporization
 Where does this occur?

On the surface of the liquid = evaporation,
throughout liquid = boiling
 Equilibrium vapor pressure depends on:


a) temperature
b) IMF’s (specific to the type of liquid)
 If a liquid has high intermolecular forces,
then what happens to that liquid’s vapor
pressure? Why?

vapor pressure ↓
high IMFs =
increase force of attraction on the particles in
the liquid. They are less likely to turn into a
gas, so vapor pressure is low.
Boiling. Freezing. Melting
 What is boiling?
 The conversion of a liquid to a vapor, within
the liquid as well as its surface. What is the
boiling point?
 The temperature at which the equilibrium
vapor pressure of the liquid is equal to the
atmospheric pressure (1 atm where we live).
 **Boiling happens throughout the
liquid…evaporation happens on the
surface.**
What is the molar heat of
vaporization?


FYI: Enthalpy is a fancy word for heat that
scientists like to use so they sound smart.
The amount of heat energy required to
vaporize one mole of liquid at its boiling point
at constant pressure.
 How does a pressure cooker work?

It elevates pressure (above 1 atm) to raise
boiling point (above 100°C) so food cooks
faster due to the higher temperature.
Freezing and melting
 What is the freezing?

The physical change of a liquid to a solid.
 What is melting?

The physical change of a solid to liquid.
 What is the molar heat (enthalpy)of fusion?

The amount of heat energy required to melt
one mole of solid at its melting point.
solid + heat  Liquid
liquid

solid + heat
re-write the equation:
solid + heat ↔ liquid
heat of fusion
Are the freezing points and melting
points the same temperature?
 Yes
 at 0°C H2O with 6kJ is a liquid
 at 0°C H2O without 6kJ is a solid
Chapter 10 Calculations
 Molar heat of Vaporization
 The amount of heat energy required to
vaporize one mole of liquid at its boiling point.
 Joules are the standard unit to measure heat
energy.
 Molar heat of vaporization for water is 40.79
kJ/mole or 2.257 kJ/gram
•
It also takes energy to melt or boil any substance. The
amount of energy required to melt or boil a substance can be
expressed by the following equations:
• q = nΔHfusion
q = change in energy (J)
n = number of moles
• q = nΔHvaporization
ΔHfusion = the molar heat of fusion
(kJ/mol)
ΔHvaporization = the molar heat of
vaporization
(kJ/mol)
•
ΔHfusion and ΔHfusion are constants and correspond to the
amount of energy it takes to freeze (fuse) or boil (vaporize)
one mole of a substance.
Ex1: How much heat energy would be required to
vaporize 5.00 moles of H2O
 q = ΔHvap·(mol)
=
=
204 kJ or 204,000 J
 Ex2: to vaporize 45.0g of H2O
 q = ΔHvap·(mol)
=
=
102 kJ or 102,000 J
 when....a liquid evaporates, it absorbs
energy. Energy is used to overcome
attractive forces. The energy doesn’t
increase the average energy of the particles,
so the temperature doesn’t change.
 when...a liquid evaporates, it takes energy
from its surroundings that’s why alcohol feels
cool to the skin.
 it’s also why we get cold when getting out of
the shower
Molar Heat of Fusion
 The amount of heat energy required to melt
one mole of a solid at its melting point.
 The molar heat of fusion of water is 6.008
kJ/mole.
 Ex1: How much energy would be required to
melt 12.75 moles of ice?

q = ΔHfus·(mol)
=
=
76.60 kJ
Ex2: to melt 6.48 x 1020 kg of ice?
 6.48x1020kg · 1000 g = 6.48x1023 g
1
1kg
 Ex3: - How much ice can be melted by 2.9 x 104 J?
 2.9x104J · 1kJ = 29kJ
1
1000J
q = (Hfus) x (mol)
29 kJ = 6.008 kJ/mol x (mol)
= 4.8 mol ice
Heat and Temperature – there is
a difference
 Heat is the amount of energy a chemical has,
frequently measured in joules (J). Because
we can’t directly measure heat, we have to
measure “temperature”, which reflects how
much kinetic energy an object has (as
measured in °C or Kelvins).
Heat and Temperature – there is a
difference
•
•
•
•
Heat transfers between objects – flows from hot to cold Law of Conservation of Energy
Ex1:ice cube in a thermos of hot water - ice melts, water
cools - same amount of heat
SI unit of heat - Joule (J) -calorie is also used frequently
•
Calorie - the amount of energy required to raise the
temperature of 1 g of water by 1 oC
• (Calories – capital letter – really means kilocalories – used
in food energy measurement)
1.000 calorie
=
4.184 Joules
This is not in your notes. Just make sure you write down the above conversion.
Three factors affect how much heat
an object absorbs or loses
 Mass of the object
 Change in temperature
 final temperature - initial temperature
 if there is no change in temperature, no heat
flows
 Specific Heat
 specific heat (Cp): heat required to raise the
temp. of 1 g of material by 1 K or 0c
 different materials have different specific heats
Specific Heat
material at
298 K and 1 atm
ice
water
steam
sodium
aluminum
iron

Cp specific heat
(J/g K)
2.09
4.18
1.86
1.23
0.9
0.45
Ex, which would you rather use to pull a pan from a hot oven- an
oven mitt or a sheet of aluminum foil? The aluminum foil will
transmit the heat easily while the oven mitt is a much better
insulator. The reason: Oven mitts have a higher heat capacity
(specific heat) than aluminum.
Computing heat to determine how much heat
is required to heat a material

It takes energy to make the temperature of
anything increase. The relationship between
energy and temperature is shown by the
equation:

q = m cp ΔT
 q = heat/energy (Joules or calories)
 m = mas - grams or moles
 Cp = specific heat
 ΔT = Tfinal - Tinitial
Specific Heat Problems
•
For water, Cp = 1.000 cal/g oC
or 4.184 J/g oC
Change your example 1 to this:
• Ex1: How much heat is required to heat 75 g of Iron (Cp = 0.444
J/gCo) from 15.5 to 57.0 oC?
•
EX2: How many calories does it take to heat 20. g of water from
10.0 to 40.0 oC? Also how many J?
Specific Heat Problems Ex3: What is the specific heat of an object if 250 calories
will heat 55 g of it from 25 to 100.0 oC?
 Ex4: - If a 100.0 g sample of silver (Cp = .237 J/g oC) at
80.0 Co loses 50. calories, what will its final temperature
be?
Temperature and Phase Changes
Flat sections at
boiling/melting why?
All energy input is
directed at changing
phase, so there is no
increase in temp.
Temperature and Phase Changes
 3 formulas to use:



q = mCpt for sections A, C, E
q = mHfus for section B
q = mHvap for section D
Temperature and Phase Change
 It is usually assumed that more heat means higher
temperature, but not when changing phase.
 EX:1 If a sample of water at 20.0 oC is heated by a hot
plate that gives off 250.0 J, how grams of water are in the
sample if the temperature rises to 30.0 oC?
Temperature and Phase Change
 EX 2: How many kJ are needed to convert
25.0 g of water from a liquid at 50.0° C to a
gas to 100.0° C.
Temperature and Phase Change
ADD TO NOTES
 EX 3: If 75 g H2O is at -5oC and is heated to 115oC.
How much total heat in joules is required? (Convert
to calories after)