CONDENSED STATES
Chemistry Spring 2015
AGENDA 3/2/15

Hard Boiled Egg??

Tests Back?

Learning Targets for Unit 3


Review from 1st Semester
Notes over the review Topics
HW: Complete the Phase Change Graphic
 **Tomorrow TESTING DAY**

ENERGY AND PHASES REVIEW
Potential vs Kinetic E
a. PE = stored energy
1.
Breaking bonds increases PE
 Making bonds releases PE
b. KE = E of motion
 KEgas > KEliquid > KEsolid

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ENERGY AND PHASES REVIEW
2. Exothermic vs Endothermic
a. Exothermic
Gives off E (E is released to surroundings)
 Bond formation
 Gas  liquid  solid (“intermolecular force
formation”)


Why? I am forming bond, making things more stable –
energy is released.
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4
ENERGY AND PHASES REVIEW
b. Endothermic
Absorbs E (E taken from surroundings)
 Bonds breaking
 Solid  liquid  gas (“intermolecular force
breaking”)


Why? Energy is needed to break bonds.
**BARF**
Break, absorb, release, form
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5
ENERGY AND PHASES REVIEW
Solid
sublimation
melting
deposition
freezing
Liquid
condensation
Gas
evaporation
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ENERGY AND PHASES REVIEW

Inside Arrows:





Exothermic (feels hot)
Decreases PE
Forming IM forces
More order
Particles closer
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ENERGY AND PHASES REVIEW

Outside Arrows





Endothermic (feels cold)
Increases PE
Breaking IM forces
Less order
Particles spread out
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ENERGY AND PHASES REVIEW
3. Heat vs Temperature
a. Heat
Form of energy
 Measured in Joules (J)  kJ (1000 J)
 Flows from warmer object to colder object

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9
ENERGY AND PHASES REVIEW
b. Temperature
Measure of average KE of particles
 How much particles are moving
 E given to particles  increase motion  increase T
 Measuring Temperature

Celsius vs Kelvin
2 different temperature scales
 K is a more direct reflection of KE

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10
ENERGY AND PHASES REVIEW
Kelvin (K)
-273oC
0K
Freezing point
of water
0 oC
273 K
100oC
373 K
Boiling point
of water
Remember: K = C + 273
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Absolute zero
3/10/2016
Celsius (oC)
11
AGENDA 3/4/15

Hand Back Unit 2 Tests


5 Minute Remediation
Notes on Solutions and Solubility

Solubility Curves

Pre-Lab for Solubility Lab Tomorrow

HW: Complete Solubility Curve WS
SOLUTIONS
Solutions
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Homogeneous  mixture is the
same throughout
 Stable  won’t spontaneously
separate
 Examples: Kool-Aid, saltwater, pop
 Made of 2 parts:
Solute – chemical that is
dissolved (Kool-Aid mix)
Solvent – chemical that does the
dissolving (water)

3/10/2016

13
SOLUTIONS


Aqueous Solution – a solution where H2O is the
solvent
Why do things dissolve?
Molecule – Ion Attraction
 Similar chemical structure

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14
SOLUTIONS
NaCl (s)
H
Na+ (aq)
+
-
+
-
-
+
-
+
+
-
+
-
O
H
H
H
O
H
O
O
H
H
H
O
+
H
O
H
H
H
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SOLUTIONS
NaCl (s)
+
-
+
-
-
+
-
+
-
+
-
Na+ (aq) +
H
H
O
H
H
O
H
H
O
-
H
H
O
O
H
3/10/2016
Cl- (aq)
H
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H
O
H
16
SOLUTIONS
 Solvation
 Occurs
on the surface of a solid
 May be exothermic or endothermic
 Heat of Solution – energy change
caused by dissolving
Ex: acid in water is warm

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AGENDA 3/5/15

Finish notes on Solutions and Solubility

Solubility Lab of Ammonium Chloride

Solubility Curves

Examples of Solubility Curves

Notes on Molarity


Practice Problems
HW: Complete Solubility Lab
WARM-UP

What is a solute?

What is a solvent?

What are three ways to dissolve solids faster?

Explain solvation to your partner.
How do you get things to dissolve faster?
SOLUTIONS

Factors that affect the rate of solvation:
1. Stirring – fresh solvent is brought to surface of
solute (faster)
2. Increase T – solvent molecules have more KE
(faster)
3. Surface area – powder dissolves faster than a
chunk
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SOLUTIONS
*LIKES DISSOLVE LIKES*
 water
& food coloring
“miscible”
2 substances are alike and will mix
water & oil
“immiscible”
2 substances are different and will not mix
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SOLUBILITY LAB

Safety:


Do not reach over the hot plate
Tips:
Heat the test tube in the water very gently
 Use the distilled water in the test tube only
 Use tap water in the beaker
 Make sure you know the amount of Ammonium
Chloride your group is responsible for.

AGENDA 3/6/15

Wrap-Up Solubility Lab


What does it mean?
Notes on Solubility Curves

Saturation/Super-Saturation

Practice Problems with Solubility

Notes on Molarity


Practice Problems
HW: Complete Solubility WS
SOLUBILITY LAB

What was the solubility of Ammonium chloride at
55 degrees C?

What would the solubility be at 85 degrees C?

What would the solubility be at 35 degrees C?

What does our best-fit line tell us about the
solubility of Ammonium chloride?
Solubility
– the maximum
amount of solute that can be
dissolved in a given amount
of solvent
SOLUBILITY CURVES
Solid in Liquid Solutions - 3 “stages”
Unsaturated
1.
•
•
Contains less than the maximum amount of solute
All solute dissolves and there is room for more
SOLUBILITY CURVES
2. Saturated
 maximum amount of solute is dissolved
 additional solute will not dissolve at that T
 depends on T
 occurs when rate of dissolving = rate of
crystallization
SOLUBILITY
3. Supersaturated
 more solute dissolved than the solution would
normally hold
 can make by carefully cooling a saturated
solution
 unstable, wants to crystallize
SOLUBILITY
= dissolving
Sugar cube
dissolving in
H2 O
= crystallizing
Some dissolves
Some crystallizes
Solution is saturated!
Equilibrium
dissolving =
crystallization
AGENDA 3/9/15

Review Solubility

Check HW

Notes on Molarity

Practice Problems with Molarity

HW: Complete Molarity WS
SOLUTIONS WARM-UP

Explain the following terms:

Unsaturated

Saturated

Supersaturated
How do we measure the amount of
solute in a solvent?
Concentration – the amount of solute
dissolved in a given amount of solvent
Dilute – small amount of solute (weak)
Concentrated – a lot of solute (strong)
MOLARITY
Molarity
 A unit of concentration
 # of moles of solute per 1 L of solution
Molarity (M) = moles of solute =
Liters of solution
mol
L
MOLARITY
Examples:
A salt solution has a volume of 152 mL and
contains 0.90 mol NaCl. What is the molarity of
the solution?
MOLARITY
M = mol
L
152 mL = 0.152 L
M = 0.90 mol
0.152 L
= 5.9 M
MOLARITY
A 3.0 L aqueous solution contains 45.0 g KI. What is
the solution’s molarity?
45.0 g KI 1 mol KI
= 0.271 mol KI
166.00 g KI
M = mol
L
M = 0.271 mol
3.0 L
M = 0.090 M
MOLARITY
How many moles of NaCl are in 1.5 L of a 3.0 M
NaCl solution? What mass of salt is this?
M = mol
L
3.0 M = x mol
1.5 L
3.0 M(1.5 L) = x mol
x = 4.5 mol
MOLARITY
4.5 mol NaCl 58.44 g NaCl
= 262.98 g NaCl
1 mol NaCl
260 g NaCl
AGENDA 3/10/15

Check Molarity HW

Warm-Up

Notes on Dilution

Salt Lab

HW: Write-Up Lab
WARM-UP

Example: A 500.00 mL solution of HCl contains
45.0 grams of the chemical. What is the molarity
of this solution?
Ans: 2.47 M
ANSWERS TO 1-5 ON SOLUTIONS PRACTICE
1.
2.
3.
4.
5.
.901M
83.9 g KCl
6.09M
1.47 x 106 g KOH
20.5 M Pb in Sn
DILUTIONS
Dilutions – making a less concentrated solution
from a stock solution
soln of known concentration
DILUTIONS
M=mol/L → M x L = mol
M1V1 = moles of solute before dilution
M2V2 = moles of solute after dilution
The moles of solute does not change during a dilution.
You’re just adding more water.
DILUTIONS
M1V1 = M2V2
DILUTIONS
Examples:
How would you prepare 305 mL of 0.500 M AgCl
from a 1.20 M stock solution?
DILUTIONS
M1 = 1.20 M
V1 = ?
M2 = 0.500 M
V2 = 305 mL = 0.305 L
M1V1 = M2V2
DILUTIONS
(1.20M)(V1) = (0.500M)(0.305L)
(1.20M)(V1) = 0.1525
1.20
1.20
V1 = 0.127 L
V1 = 127 mL
127 mL of stock soln
& 178 mL of H2O
DILUTIONS
To prepare 1.00 L of 6.0 M HCl, what volume of
concentrated (12.1 M) HCl must be used?
M1 = 12.1 M
V1 = ?
M2 = 6.0 M
V2 = 1.00 L
M1V1 = M2V2
DILUTIONS
(12.1 M)(V1) = (6.0 M)(1.00 L)
(12.1 M)(V1) = 6.0
12.1
12.1
V1 = 0.496 L
V1 = 0.50 L
V1 = 500 mL
500 mL of 12.1 M HCl
should be added
to 500 mL of H2O
SALT SOLUTIONS LAB

Safety:


Bunsen Burners: Keep your hair tied back
Tips:

Have the fresh samples of water ready to go while
boiling the first solution and salt solutions.

Keep data on the back of the lab for easy reference.

Be sure to use distilled water, not tap water.
AGENDA 3/11/15

1st Hour: Complete Lab, graph and conclusion

7th and 8th:


Review Findings in the lab

Practice Dilution and Molarity calculations
HW: Finish any make-up work you have left
SALT LAB WARM-UP

Answer these questions with your partner:
1.
How did adding the salt to the water affect the
boiling temperature of the water?
2.
How did diluting the salt concentration affect the
boiling temperature of the water?
3.
Bonus question: Why would adding salt to pasta
water make the pasta cook faster?
CONCLUSION



Take time to write up your lab conclusion.
When you are complete turn in your lab and pick
up the dilutions worksheet.
When finished with the dilutions worksheet, pick
up the molarity and dilution worksheet.
AGENDA 3/12/15

1st Hour: Collect Lab

Check Dilutions Hw

Notes on Evaporation and Boiling

Boiling Demonstration

Colligative Properties

HW: Complete WS as study guide for Quiz
SALT LAB WARM-UP

Answer these questions with your partner:
1.
How did adding the salt to the water affect the
boiling temperature of the water?
2.
How did diluting the salt concentration affect the
boiling temperature of the water?
3.
Bonus question: Why would adding salt to pasta
water make the pasta cook faster?
PHASE CHANGE: EVAPORATION
Evaporation –
 Process of changing from l  g
 Occurs at any T
 Occurs at the surface of the
liquid
PHASE CHANGE: EVAPORATION
Saturated air with vapor (“air is
full”)
PHASE CHANGE: EVAPORATION
Vapor Pressure
The pressure (collisions) caused by the vapor above
a liquid in a sealed container of equilibrium
Caused by the collision of vapor
molecules with the surface of the
liquid
PHASE CHANGE: EVAPORATION
VP depends on:
 T (as T ↑, evaporation ↑, vapor ↑ and VP ↑)


Because more molecules have the E to break free of the
surface
T (as T ↓, evaporation ↓, vapor ↓ and VP ↓)

Because fewer molecules have E to break free of the
surface
PHASE CHANGES: BOILING
Occurs when the VP of a liquid = the atm P on it
 Vapor forming within the liquid, not at surface
 The pressure from the atmosphere would crush
bubbles

P inside =
P outside
PHASE CHANGE: BOILING
Boiling Point –
 T at which VP = atmospheric P
 Normal Boiling Point (NBP) = BP when P = 1 atm
P inside =
P outside
PHASE CHANGE: BOILING


Why would the vapor pressure need to be equal to
or greater than the atmospheric pressure?
How could we increase the vapor pressure?


1 atm
Think how to increase the amount of vapor present in a
liquid.
ANSWER: We heat it…vapor pressure is directly
related to the temperature.
.022 atm
PHASE CHANGES

What if you’re not at 1 atm?
Pike’s Peak, CO
CLC, IL
PHASE CHANGES
1.05 atm (mine)
107oC
1.00 atm (IL)
100oC
0.894 atm (Denver)
94oC
Pike’s Peak
89oC
Mt. Everest
70oC
H2O VP = 0.022 atm
PHASE CHANGES


Will foods cook faster in Denver?
No, because the BP is lower, water is cooler,
cooking takes longer.

What causes the boiling point and melting point
to change when a salt is dissolved in it?
COLLIGATIVE PROPERTIES
Properties that depend
entirely on the number of
particles dissolved in a
given mass of solvent
Ex: Boiling point and melting point
COLLIGATIVE PROPERTIES
Vapor Pressure Lowering



Allows liquids to boil at temperatures lower than
their normal boiling point.
Vacuum Demo!!
COLLIGATIVE PROPERTIES
Why?
= solvent
= solute
Because vaporization occurs at the surface and
the solute is in the way….so less vaporizes,
therefore VP is lower.
*depends on the # of particles of solute
COLLIGATIVE PROPERTIES
2. Boiling Point Elevation – a solution boils at a higher T
than a pure solvent (like our salt lab)
Why?
 Boiling occurs when VP = atm P
 Since VP is lowered by adding salt, more E must be
added to raise VP to atm P
 Proportional to # of particles
COLLIGATIVE PROPERTIES
Example:
Salt in H2O for
cooking pasta
 H2O
boils at a
higher T, pasta
cooks faster
 H2O does not boil
faster
COLLIGATIVE PROPERTIES
3. Freezing Point Depression – a solution freezes
at a lower T than a pure solvent
Why?
 Particles in liquid disrupt the orderly pattern of
solid crystal, bonds cannot form as easily
 Must get colder to solidify
 Proportional to the # of particles
COLLIGATIVE PROPERTIES
Example:
Melting Ice
Sand doesn’t melt ice
 NaCl dissolves, melts ice by
lowering fp
 AlCl3 better (more particles)

AGENDA 3/13/15

Check Molarity and Dilutions HW

Last minute questions before quiz

Molarity and Dilutions Quiz
AGENDA 3/16/15

Quizzes Tomorrow/Maybe Today

Back to Balancing Equations

Equations in Solutions
Predicting States of matter
*** GRAB A WS FROM THE BACK TABLE***

HW: Ionic Equations
SOLUTIONS
NaCl (s)
H
Na+ (aq)
+
-
+
-
-
+
-
+
+
-
+
-
O
H
H
H
O
H
O
O
H
H
H
O
+
H
O
H
H
H
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AGENDA 3/17/15 ST. PATTY’S DAY!!

Review Quiz

Ionic Equation Review


HW Check
Ionic Reaction Demo
Energy/Phases Review

HW: Energy and Ionic Equation Worksheet
IONIC REACTIONS DEMO

I have a solution of lead (II) nitrate and lithium
chloride.

Will a precipitate form?

If not, why not?

If yes, what is the precipitate that formed?

Write a complete balanced reaction including phases.
IONIC REACTION DEMO

I have a solution of sodium hydroxide and a
solution of hydrogen chloride.

Will a precipitate form?

If not, why not?

If yes, what is the precipitate that formed?

Write a complete balanced equation including phases.
PHASE DIAGRAMS
Describe the conditions of T and P where each
phase is most likely to exist
 Gas



Liquid


High T, low P
middle
Solid

Low T, high P
82
PHASE DIAGRAMS
 Phase
Diagram – shows the
relationship between the 3 phases in
a sealed container
83
PHASE DIAGRAMS
Features:
Lines of Equilibrium




1.
2.
3.
On the lines, the substance exists in equilibrium between
two states
At that T and P, both phases exist
s
s
l
g
l
g
84
PHASE DIAGRAMS
Features:
 Triple Point

One specific point at which all three phases exist in
equilibrium in a sealed container
 A

85
PHASE DIAGRAMS
 Features:
 Critical
Point
 Point
past which neither liquid nor gas
truly exist
 An intermediate fluid phase (plasma*)
B
86
PHASE DIAGRAMS
87
Phase Diagram
for H2O