Download Liquids

Chemistry
Chapter 13 Phases
These notes are set up as a Guided Reading Packet. Use your textbook to
complete the assigned sections of the notes as you read the material.
Liquids
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, the motion of liquid particles can be described as….
Properties of Liquids and the Particle Model – define each property:
1. Definite volume
2. Fluidity
3. Relative high density
4. Incompressible
5. Dissolving ability
6. Ability to diffuse
7. Surface tension
8. Tendency to evaporate and boil
9. Tendency to solidify
Questions:
Notes_S
Page 1 of 13
Chemistry
Chapter 13 Phases
1. Why are liquids more dense than gases?
2. Why are liquids harder to compress than gases?
3. Why do liquids diffuse slower than gases?
4. Can a liquid boil without increasing the temperature? How?
Solids
“Solid as a rock,” is the intuitive 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 ______________________, less __________________, more packed ______________________,
and higher _________________________________
Properties of Liquids and the Particle Model – define each property:
Notes_S
Page 2 of 13
Chemistry
Chapter 13 Phases
1. Definite shape and volume
2. Non-fluid
3. Definite melting point
4. High density
5. Incompressible
6. Slow diffusion
Crystalline Solids
Classification of crystals by arrangement and shape
Crystal Lattice (define) ____________________________________________________________________________________
____________________________________________________________________________________
____________________________________________________________________________________
____________________________________________________________________________________
The smallest portion of the crystal lattice that reveals the 3-D pattern of the entire lattices is the
____________________.
Binding Forces in Crystals
Simple
Notes_S
Body-centered (ex. Li,
K, Cr)
Face-centered (ex. Cu,
Ag, Au)
Hexogonal (like oranges
in a grocery store); (ex.
Zn)
Page 3 of 13
Chemistry
Chapter 13 Phases
Binding forces in crystals – complete the table from reading in your book
1.
Binding Force
Lattice consists of
Formed When / Binding Force
1. Ionic crystals
Group 1/2 metals combine with
Group 7/8 nonmetals
2. Covalent network crystals
Atoms bond to neighbors,
extending through a network,
large chains form
3. Metallic crystals
Each e- and the (+) metallic ions
attract electrostatically
4. Covalent molecular crystals
For nonpolar molecules, London
Forces; For polar molecules,
Dipole-Dipole.
Amorphous Solids
Rubber, glass, plastics and synthetic fibers are called amorphous solids.
“Amorphous,” comes from the Greek for “________________________.”
Unlike crystals, amorphous solids do not have a regular, natural shape, but instead take on whatever
shape imposed on them.
Particle arrangement is not _____________________________________________________________.
Examples of amorphous solids –
____________________________________________________________________________________
____________________________________________________________________________________
Amorphous solids are prepared by ________________________________________________________.
Notes_S
Page 4 of 13
Chemistry
Chapter 13 Phases
Molecular examples
Changes of State
Complete the table as on page 342
Possible Changes of State
Change of State
Name
Example
Equilibrium
1. What does equilibrium mean?
_________________________________________________________________________________
_________________________________________________________________________________
_________________________________________________________________________________.
1. What is a closed system?
Draw figure 13:
Notes_S
Page 5 of 13
Chemistry
Chapter 13 Phases
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 –
______________ + ______________
_____________
_____________ + _______________
_____________
When a vapor condenses, as in condensation, it gives off heat energy and can be shown as:
And condensation –
___________
_____________ + ________________
The liquid vapor equilibrium can be rewritten as:
________________ + _________________
__________________
“The double yields sign represents a reaction at _____________________________”
Le Chatelier’s Principle
1. What is it?
2. Is temperature an example of stress?
3. What happens when you increase the temperature of a system?
4. What happens when you decrease the temperature of a system?
5. What factor is controlling the decrease and increase of vapor and liquid?
Notes_S
Page 6 of 13
Chemistry
Chapter 13 Phases
Equilibrium Vapor Pressure of a Liquid
1. What is it?
______________________________________________________________________________
______________________________________________________________________________
_____________________________________________________________________________.
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.
At 80° C
At 50° C
At 20° C
Water
355 torr
92 torr
20 torr
Alcohol
760 torr
400 torr
90 torr
Cooking Oil
10 torr
4 torr
1 torr
2. What is equilibrium called when liquid molecules enter into the gaseous state?
______________________________
Notes_S
Page 7 of 13
Chemistry
Chapter 13 Phases
3. Where does this occur?
______________________________________________________________________________
______________________________________________________________________________
4. Equilibrium vapor pressure depends on:
a) ________________________________________________________________
b) ________________________________________________________________
5. If a liquid has high intermolecular forces, then what happens to that liquid’s vapor pressure? Why?
vapor pressure ↓
high IMFs = ________________________________________
Boiling. Freezing. Melting
Boiling
1. What is boiling?
2. What is the boiling point?
3. What is the molar heat of vaporization?
4. How does a pressure cooker work?
Freezing and melting
1. What is the freezing?
2. What is melting?
3. What is the molar heat of fusion?
Notes_S
Page 8 of 13
Chemistry
Chapter 13 Phases
solid + heat
_____________________
re-write the equation:
solid + heat
liquid
heat of fusion
4. Are the freezing points and melting points the same temperature?
at 0°C H2O with 6kJ is a liquid
at 0°C H2O without 6kJ is a solid
Notes_S
Page 9 of 13
Chemistry
Chapter 13 Phases
Calculations
Heat of Vaporization
The amount of heat energy required to vaporize one unit (mass or moles) of liquid
q = (H vap ) x ( unit )
Joules are the standard unit to measure heat energy.
unit = gram or mole
Molar heat of vaporization for water is 40.79 kJ/mole or 2.257KJ/g
Ex1: How much heat energy would be required to vaporize 5.00 moles of H2O
Ex2: to vaporize 45.0g of H2O
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
Heat of Fusion
The amount of heat energy required to melt one unit (mole or gram) of a solid at its melting point.
The molar heat of fusion of water is 6.008 kJ/mole or 0.334KJ/gram
q = (Hfus) x (unit)
Ex1: How much energy would be required to melt 12.75 moles of ice
unit = g or mole
Ex2: to melt 6.48 x 1020 kg of ice
Ex3: - How much ice can be melted by 2.9 x 104 J?
Notes_S
Page 10 of 13
Chemistry
Chapter 13 Phases
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). In thermodynamics, the term “enthalpy” is used interchangeably with
“heat”, as it avoids confusion between the terms “heat” and “temperature”.
•
•
•
•
Heat transfers between objects – flows from hot to cold - Law of Conservation of Energy
•
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)
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
1 calorie = 4.184 Joules

Three factors affect how much heat an object absorbs or loses
o
o

final temperature - initial temperature

if there is no change in temperature, no heat flows

specific heat (Cp): heat required to raise the temp. of 1 g of material by 1 K

different materials have different specific heats
o
material at
Cp specific heat
298 K and 1 atm
(J/g K)
o
Notes_S
ice
2.09
water
4.18
steam
1.86
sodium
1.23
aluminum
0.9
iron
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.
Page 11 of 13
Chemistry
Chapter 13 Phases

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 =
Ex1: How much heat is required to heat 75 g of Iron (Cp = 0.444 J/gCo) from 15.5 to 57.0 oC?
EX 2: How many joules does it take to heat 20. g of water from 10.0 to 40.0 oC? Also how
many calories?
EX 3: 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?
Notes_S
Page 12 of 13
Chemistry
Chapter 13 Phases
Temperature and Phase Changes
flat sections at boiling/melting
why? all energy input is directed
at changing phase, so there is no
increase in temperature
3 formulas to use:
q = mCpt for sections ______________________
q = mHfus for section _______
q = mHvap for section _______
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?
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.
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)
Notes_S
Page 13 of 13