Matter & Energy
Honors Chemistry
Science
Science is a body of knowledge collected
by scientists over many years & the
methods used to obtain the knowledge
B. Chemistry is the study of the composition,
structure and properties of matter & the
changes it undergoes
A.
 Chemical = any substance that has a definite
composition
1.
2.
It is through the analysis of much information
on matter that we can solve problems & answer
question
What, how much, how it can be changed, &
how fast
States of Matter
Solid
condensing
vaporizing
Liquid
Gas
Vaporization

Boiling – conversion of a liquid to a gas within the
liquid as well as at its surface


Boiling Point – vapor pressure of the liquid equals the
atmospheric pressure
Evaporation – particles escape from the surface of a
non-boiling liquid and enter the gas state

Particles at the surface have higher than avg energies that
overcome the intermolecular forces that bind them to the
liquid
Holt Visual
Vaporization and
Condensation
State
Solid
Shape
Definite
Volume
Definite
Liquid Indefinite Definite
Gas
Indefinite Indefinite
Movement
Structure
Particles only
vibrate about
fixed points
Particles packed
together in
relatively fixed
positions;
strong attractive
forces
Particles can
move past
one another
Particles move
more rapidly –
temporarily
overcome strong
attractive forces;
allows flow
Particles
move very
rapidly
Particles are at
a great distance
from each other;
attractive forces
weak
Intermolecular Forces
 Liquid

Nitrogen
Boils at -196C
 Mercury
Liquid at room temp
 Freezes at -39C

 Transfer
of Heat from Hg to N2
Properties
1.
Physical property - can be observed
without changing the identity of the
substance
•
Intensive is independent of amount

•
Extensive is dependent of amount

2.
mp, bp, density, conducts electricity/heat, temp
mass, volume, amount of energy, heat
Chemical property – relates to a
substance’s ability to undergo
changes that transform it into
different substances
Changes
Physical – does not involve a change
in the identity of a substance; may
change the appearance
Chemical – one or more substances
are converted into different
substances with different properties
1.
2.


Alters identity of substance. Produces a
new substance
The new substance (product) has
different properties than the beginning
materials (reactants).
Signs of a Chemical
Change
Color
Gas (change in odor)
Precipitate
Change in temperature (may include
light)
1.
2.
3.
4.


Endothermic vs. Exothermic reactions
Note: all chemical and physical changes
involve energy
What is the 3rd change?
Nuclear Change - changes the
composition of the atom’s nucleus
3.




tremendous amount of energy involved
Fission vs. Fusion
Radioactive decay
Where is uranium?
 Ground
 Refined for nuclear power plants
Radioactive Decay
Conservation Matter
and Energy

Cannot be created or destroyed, only
changes form in a chemical or physical
change
 Burning magnesium
• Burn Mg – heavier product, why?
• Mg + O2  MgO

Types of Energy: electrical, mechanical,
light, chemical mechanical, thermal
• Heater – electrical energy to heat energy
• Photosynthesis – light to chemical
• Transportation – chemical to thermal to
mechanical
E. Classification
MATTER
Anything that has
mass and volume
Pure Substances
Mixtures
Fixed composition;
characteristic chemical
& phys properties
Blend of 2/more kinds
of matter, each of
which retains its own
identity & properties
Elements
Compounds
Periodic table;
smallest particle
to retain all
properties atom
2/more different
elements
chemically
bonded (I or C)
H2O vs H2O2
Homogeneous
(Solution)
Uniform in
composition same proportion
of components
throughout
Heterogeneous
Not uniform
throughout
The Periodic Table
Metals









Location: to the left of the staircase
At room temp, all are solid except for Hg
Ductile - can be drawn out into thin wires
Malleable - can be hammered into thin
sheets
Luster (A.K.A. Shininess)
Good conductors of heat and electricity
High density
High melting points
Ion formation – tend to lose electrons
resulting in positive charges
Nonmetals
 Location:
to the right of the staircase
 At room temperature, they are solids,
liquids, or gases.
 Dull – no luster
 Insulators of heat and electricity.
 Brittle - Neither malleable or ductile
 Lower bp and mp than metals.
 Ion formation – tend to gain electrons
resulting in negative charges
Metalloids (Semimetals)

Located between the metals and
nonmetals, ALONG the staircase .

Have properties of both metals and
nonmetals.

There are 7 metalloids in the periodic
table: Boron (B), Silicon (Si), Germanium
(Ge), Arsenic (As), Antimony (Sb), Tellurium
(Te), & Astatine (At).
Check for Understanding
1.
List the nonmetals in the 5th period.

2.
Metalloid(s) in group 5A (15)?

3.
Arsenic and Antimony
Liquid metal? Liquid nonmetal?

4.
Iodine and Xenon
Mercury
Bromine
Symbol for the ion in group 6A and
period 3?

S-2
Compounds
2
or more elements chemically
combined through covalent or ionic
bonding
 Examples:





Na and Cl2 react to form NaCl
C and O2 react to form - CO2
How many atoms in NH4Cl?
How many H atoms (NH4)2SO4
How many H atoms in 5 (NH4)2SO4
Solid Solutions
Alloys: Solid solutions containing two or more
metals or a metal and a nonmetal
 Advantages of alloys over pure metals:
 Stronger, cheaper, resistant to corrosion, lighter,
harder

Brass is an alloy of
copper and zinc.
Steel is an alloy of
carbon and iron.
Stainless steel
contains chromium
Bronze is an alloy
of copper and tin.
A closer look at alloys
Alloy
Metals
Yellow Gold
(14 or 18 carat)
Gold, Silver, Copper
Red Gold
Gold and Copper
White Gold
Gold and Palladium
Sterling Silver
Silver and Copper




Suspensions are
mixtures of particles that
settle out if let undisturbed.
Heterogeneous
Suspensions can be filtered,
while solutions cannot.
Blood, aerosols, OJ





Colloids are a type of
mixture whose particles are
held together through
Brownian Motion, the
erratic movement of colloid
particles.
Colloids cause the Tyndall
Effect, or scattered light
due to Brownian motion.
Intermediate between
homogeneous and
heterogeneous
The size of the particles is
smaller than those found in
suspensions and greater
than those found in
solutions.
Milk, paint, fog , smoke,
dust
Colloids
Tyndall effect is caused by reflection of light by very
small particles in suspension in a transparent medium. It
is often seen from the dust in the air when sunlight
comes in through a window, or when headlight beams are
visible on foggy nights
Shows the scattering of light by shining lasers of different colors
through colloids and water.
The laser beam is visible through the colloid.
Separation Techniques
 Heterogeneous



Mixtures
Filtration: Pour liquid through filter
paper to collect solid
Centrifuge: separates solid-liquid
mixtures
Decanting
Separation Techniques
 Homogeneous



Mixtures
Crystallization: evaporate liquid and
solid will crystallize
Chromatography – used to separate
pigments of ink on a strip of paper.
Distillation
Distillation - separation of a solution
based on differences in boiling point
Compounds

Decomposition – compound breaks down into
two or more simpler compounds or elements
Electrolysis - decomposes a
compound with electricity
% Concentration of
Solutions
 Solute
 Solvent
 Solution
Solute x 100% = % Concentration
Solution
Saturated – soln
containing the max
amt of solute
 Unsaturated – soln
containing less
solute than a sat
soln under the
existing conditions
 Supersaturated –
contains more
dissolved solute
than a saturated
solution under the
same conditions

Solubility Curves
supersaturated solution
(stirred)
Solubility
(physical change)

Definition: mass of
solute needed to make
a saturated solution at
a given temperature



solution equilibrium in
a closed system
dissolution ↔
crystallization
Unit = g solute/100 g
H2O
At 20oC, a saturated
solution contains how
many grams of NaNO3
in 100 g of water?
90 g
What kind of solution is
formed when 90 g
NaNO3 is dissolved in
100 g water at 30oC?
unsaturated
What kind of solution
is formed when 120 g
NaNO3 is dissolved in
100 g water at 40oC?
supersaturated
180
Saturated sol’n
170
160
150
140
Supersaturated
solution
130
120
Solubility ( g/100 g water )
What is the solubility
at 70oC?
135 g/100 g water
Solubility Graph for NaNO3
110
100
90
80
70
Unsaturated solution
60
50
40
30
20
10
0
0
10
20
30
40
50
60
70
Temperature (deg C)
80
90
100
110
Solubility of solids in liquids
For most solids, increasing
temperature, increases solubility.
 In general, “like dissolves like”.
Depends on




Type of bonding
Polarity of molecule
Intermolecular forces between solute and
solvent
Solubility of Gases
Gases are less
soluble at high
temperatures than
at low temperatures
 Increasing
temperature,
decreases
solubility.
 Increasing
pressure, increases
solubility.


The quantity of gas that dissolves in a
certain volume of liquid is directly
proportional to the pressure of the gas
(above the solution).

Effervescence – rapid escape of gas
dissolved in liquid
Factors Affecting Solubility
Increase surface area of solute
(crushing)
 Stir/shake
 Increase temperature

Energy Concepts


Thermochemistry: the study of the changes in
energy that accompany a chemical reaction and
physical changes.
Chemical Reactions involve changes in energy
that result from
•
•

Bond breaking that requires energy (absorbs) from the
surroundings.
Bond making that produces energy (releases) to the
surroundings.
Changes in energy result in an energy flow or
transfer.
Heat vs. Temperature

Heat: (q) is the energy
transferred due to changes
in temperature.
•

Temperature (T) is a measure
of the average particle motion
or the average kinetic energy.
• Heat flows spontaneously
from a higher to a lower
temperature.
Heat vs. Temp Simulation - Eureka
Calorimeter
Heat is measured
in a calorimeter.
Changes in
temperature are
measured in a
known quantity of
water in an
insulated vessel.
Simple calorimeter used
in class = Styrofoam cup
Types of Reactions
Exothermic: releases heat into their
surroundings.
1.


Heat is a product and temperature of the
surroundings increase.
This occurs during bond formation.
surroundings
surroundings
Exothermic
Reaction
(system)
surroundings
surroundings
Types of Reactions
Endothermic: absorbs heat from the
surroundings.
2.


Heat acts as a reactant and temperature of the
surroundings decreases.
This occurs during bond breaking.
surroundings
surroundings
Endothermic
Reaction
(system)
surroundings
surroundings

Exothermic Example:
Dissolving calcium chloride in water
CaCl2

(s)
H2 O

 Ca +2
(aq)
+ 2Cl-1
(aq)
+ 88.0kJ
Combustion reactions are ALWAYS exothermic:
C3H8 (g) + 5O2 (g) → 3CO2 (g) + 4H2O(g) + 2043 kJ

Endothermic Example:
2NH4Cl (s) + Ba(OH)2·8H2O (s) + 63.9 kJ 
BaCl2 (s) + 2NH3 (g) + 10H2O (l)

Physical states are written – influences the overall energy exchanged.
Very specific!
Forms of Energy
Mechanical, Heat, Chemical, Electrical,
Radiant, Sound, Nuclear
Changes of State
A.
Energy
1.
Types
a) Potential energy is the energy of position
1)
As particles move apart, the PE increases
2)
The PE of a gas is greater than the PE of
a liquid which in turn is greater than the
PE of a solid
3)
During condensation, the PE decreases
and energy is released. This is an
exothermic change.
Changes of State
Kinetic energy is the energy of motion.
b)
1)
2)
3)
4)
Except at 0 K, all particles are in constant
motion
Temperature is a measure of the avg KE of the
particles in a sample.
When temperature is increased, the KE of the
particles increases.
In a liquid, the particles must have a minimum
KE (Em) in order to overcome the
intermolecular attractions of neighboring
particles to escape.
•
The stronger the intermolecular forces in a liquid,
the higher the Em.
Heating and Cooling Curves
graph of temp of a substance
Label the Heating Curve of Water
2. Evaluate the energy changes that occur
during a heating curve.
1.
•
•
•
•
Hf – heat of fusion: energy needed to melt an
amount of a substance at its mp
Hv – heat of vaporization: energy needed to
vaporize an amount of a substance at its bp
Hf and Hv Units: J/g or kJ/mol or cal/g
Hf and Hv are physical properties of a
substance
Heating Curve for Water
Temperature (ºC)
Hv =2259 J/g
q= mCΔT
lg
q= mHv
D
C = 1.841 J/g°C
E
100
q= mCΔT
Hf = 334 J/g
B
0
sl
q= mHf
q= mCΔT
C
C = 2.092 J/g°C
A
Energy
C = 4.184 J/g°C
Problems
1.
Calculate the energy (in cal) needed to
melt 125.0 g of ice at 0.0°C
9,978 cal
2.
How much energy (in kJ) is needed to
warm 180.0g of ice at -20.0°C to water at
75.0°C?
124.2 kJ
3.
If 275.0 g of liquid water at 100.0°C and
475.0 g at 30.0°C of water are mixed in an
insulated container, what is the final
temperature?
55.7°C
Physical Properties of Gases:
1. Gases consist of small particles that have
mass. These particles are usually molecules,
except for the noble gases.
Physical Properties of Gases:
2.
Gases have mass. The density is
much smaller than solids or liquids,
but they have mass. (A full balloon
weighs more than an empty one.)
3.
The particles in gases are separated
by relatively large distances. Gases
can be compressed. It is very easy
to reduce the volume of a gas.
Unlike liquids,
gases completely
fill their
containers.
5. The particles in
gases are in
constant rapid
motion (random).
4.
6.
Gases can move through each other
rapidly - diffusion (ex. food smells and
perfume)
7. Gases exert
pressure because
their particles
frequently collide
with the walls of
their container and
each other.
8. Collisions of gas particles are elastic.
Inelastic Collision
Elastic Collision
Gas particles do not slow down when hitting
each other or the walls of their container.
9. Gas particles exert no force on one another.
Attractive forces are so weak between particles
they are assumed to be zero.
10. Temperature of a gas
is simply a measure of
the average kinetic
energy of the gas
particles.
High temp. = high KE
Low temp. = low KE
The pressure of a gas depends upon
temperature
high temp. = more collisions, high pressure
low temp. = less collisions, low pressure
Low pressure
High pressure
Boyle’s Law
Pressure - Volume Relationship
 The
pressure & volume of a sample of gas at
constant temperature are inversely proportional
to each other. Law assumes n (amount) is
constant.
Inverse
P1V1 = P2V2
Boyle’s Law
V ____P ____
more collisions
(smaller volume, ____________)
Boyle’s Law Problem
A sample of oxygen occupies 300. mL under a
pressure of 740. mm Hg. If the temperature
remains constant, calculate the volume under a
pressure of 750. mmHg.?
V1 = 300. mL
P1 = 740. mm Hg
V2 = ?
P2 = 750. mm Hg
V2 = 296. mL
Charles’ Law:
Temperature - Volume Relationship.
At constant pressure the volume of a fixed
amount of gas is directly proportional to its
absolute temperature. Law assumes n is
constant.
Direct
V1
V2
=
T1
T2
*Temperatures must be in Kelvin!
K = °C + 273
Balloon in cool and cold water:
Charles’s Law
Charles’s Law Problem
A
gas sample at 83ºC occupied a
volume of 1470 m3. At what
temperature, in ºC, will it occupy a
volume of 1250 m3?
V1 = 1470 m3
T1 = 83°C = 356 K
V2 = 1250 m3
T2 = ?
T2 = 30.°C
Gay-Lussac’s Law
Pressure-Temperature Relationship
 The
pressure of a fixed volume of gas is
directly proportional to its absolute
temperature. Law assumes n is constant.
Direct
P1 = P2
T1
T2
*Temperatures must be in Kelvin!
K = °C + 273
Gay-Lussac’s Law
T ____ P ____ (moves faster,) more collisions
Gay-Lussac’s Law Problem
Before a trip, the pressure in a car tire was 1.80
atm at 21oC. At the end of the trip, the pressure
gauge reads 1.90 atm. Calculate the
temperature, in Celsius, of the air inside the tire
at the end of the trip. Assume the tire volume
does not change.
P1 = 1.80 atm
P2 = 1.90 atm
T1 = 21°C = 294 K
T2 = ?
T2 = 37°C
The Combined Gas Law (“Choyles”)
Pressure-Volume-Temperature relationship
This law can be used to determine how changing
two variables at a time affects a third variable.
P1V1 P2V2
=
T1
T2
Combined Gas Law Example: A gas occupies 72.0 mL
at 25 °C and 198 kPa. Convert these to standard
conditions. What is the new volume?
P1 = 198 kPa
P2 = 101.325 kPa
P1V1
P2 V2
V1 = 72.0 mL
V2 = ?
=
T1
T2
T1 = 298 K
T2 = 273 K
198 kPa  72.0 mL  = 101.325 kPa  V2
298 K
129 mL = V2
273 K
Dalton’s Law of
Partial Pressure
 Gases
in a mixture behave
independently of each other.
 The total pressure of a gaseous mixture
equals the sum of the partial pressures of
the individual gases in a mixture.
 Partial pressure = individual
pressure of a gas in a mixture
PT = p1 + p2 + p3 + …
Dalton’s Law of Partial Pressures:
PT = Pa + Pb + Pc + …
Example #1) A flask contains a mixture of oxygen, argon, and
carbon dioxide with partial pressures of 745 torr, 0.278 atm,
and 391 torr respectively. What is the total pressure in the
flask?
 760 torr 
.278 atm 
 = 211 torr
 1 atm 
+ 745 torr
+ 391 torr
1347 torr
Dalton’s Law of
Partial Pressure
 In
the lab, gases are collected over water
(water displacement). As a result, water
vapor contributes to the total pressure.
PT = pdry gas + pwater vapor
where pwater vapor varies with temperature
T (oC)
P (mm Hg)
T (oC)
P (mm Hg)
T (oC)
P (mm Hg)
T (oC)
P (mm Hg)
0
4.6
26
25.2
51
97.2
76
301.4
1
4.9
27
26.7
52
102.1
77
314.1
2
5.3
28
28.4
53
107.2
78
327.3
3
5.7
29
30.0
54
112.5
79
341.0
4
6.1
30
31.8
55
118.0
80
355.1
5
6.5
31
33.7
56
123.8
81
369.7
6
7.0
32
35.7
57
129.8
82
384.9
7
7.5
33
37.7
58
136.1
83
400.6
8
8.1
34
39.9
59
142.6
84
416.8
9
8.6
35
42.2
60
149.4
85
433.6
10
9.2
36
44.6
61
156.4
86
450.9
11
9.8
37
47.1
62
163.8
87
468.7
12
10.5
38
49.7
63
171.4
88
487.1
13
11.2
39
52.4
64
179.3
89
506.1
14
12.0
40
55.3
65
187.5
90
525.8
15
12.8
41
58.3
66
196.1
91
546.1
16
13.6
42
61.5
67
205.0
92
567.0
17
14.5
43
64.8
68
214.2
93
588.6
18
15.5
44
68.3
69
223.7
94
611.0
19
16.5
45
71.9
70
233.7
95
634.0
20
17.5
46
75.7
71
243.9
96
658.0
21
18.7
47
79.6
72
254.6
97
682.0
22
19.8
48
83.7
73
265.7
98
707.3
23
21.1
49
88.0
74
277.2
99
733.2
24
22.4
50
92.5
75
289.1
100
760.0
25
23.8
Eudiometer

Piece of glassware used to
measure the change in
volume of a gas. It is
similar to a graduated
cylinder. It is closed at the
top end with the bottom end
immersed in water or
mercury. The liquid traps a
sample of gas in the
cylinder, and the graduation
allows the volume of the gas
to be measured.
Example #2) Atmospheric pressure is 101.3kPa,
and air is a mixture of N2, O2, and Ar as 78.0%,
21.0%, and 1.0%, respectively. Calculate the
partial pressure of O2.
21.3 kPa
Example #3) Hydrogen gas is collected by water
displacement at 18°C. Air pressure on that day
is 744.0 mm. Calculate the pressure due to the
dry hydrogen gas.
728.5 mm Hg