Matter and Energy
Honors Chemistry
MATTER is anything that has mass and
volume.
Do you
know ALL
the states
of matter?
Kinetic Molecular Theory of Matter
• Assumptions:
• All matter consists of particles, such as, atoms, molecules,
formula units.
• Particles are in constant motion (kinetic). These motions
are
•
vibrational (only for solids),
• translational, and/or
• rotational.
All 3 for liquids
and gases.
• Collisions are elastic.
• Theory explains the physical properties of matter.
• The state of matter of a substance depends upon the
strength of attraction between particles.
Vibrational
Rotational
Translational
Combined Motions
Gases
gas
Liquids
liquid
Solids
solid
low strength of attraction
high strength of attraction
little to no order
most order
high density (condensed states)
low density
incompressible
compressible
no definite volume or shape
definite volume
takes shape of container
diffuses quickly
slower diffusion ( with  temp)
definite volume
definite shape
low diffusion
States of Matter
Solid
condensing
evaporating
Liquid
Gas
Properties of Matter
 Physical = A characteristic of a substance that does
not involve a chemical change
 Examples: texture, state of matter, density, hardness,
boiling point
Properties of Matter
 Chemical = A property of matter that describes a
substance’s ability to participate in chemical
reactions.
 Examples: reacting with oxygen, light sensitivity
Physical Changes
 Do NOT change the identity
 Often change what the substance looks like
 Examples: mixing ice tea in water, crushing a
rock, freezing water
Chemical Changes
 Alter the identity of the substance.
 The new substance has a different composition than
the beginning substances.
 Examples: rusting and burning
 A shorthand way to express a chemical reaction is
with a chemical equation.
Chemical Equations
A+B
C+D
 The substances on the left side of the arrow are called the reactants.
They are the starting materials in the reaction.
 The substances on the right side of the arrow are called the products.
They are the ending materials in a reaction.
Examples of Chemical Equations
2 H 2O  2 H 2  O2
H 2O2  H 2O  O2
N 2  3H 2  2 NH 3
Signs of a Chemical Change
 Gas production – bubbles, odor, fizz, smoke
 Color change
 Release or absorption of energy – light or
temperature change
 Formation of a precipitate – a solid substance that
falls out of solution
Classification of Matter
Classification of Matter
 Pure substances: A sample of matter with a
definite composition; means definite chemical
and physical properties.
 Includes: Elements and Compounds
Elements
 Made up of one type of atom.
 An atom is defined as the smallest unit of an
element that maintains the properties of that
element.
 Cannot be separated into simpler substances by
chemical means.
 Represented by symbols.
Elements (cont.)
 Can exist as atoms or molecules. A molecule
usually consists of two or more atoms.
 Ex. N2, O2, F2, Cl2, Br2, I2, H2, P4, S8
 Elements that have more than one form are
called allotropes.
 Ex. Carbon (graphite and diamond)
Parts of the Periodic Table
Metal, Nonmetals and Metalloids (Semimetals):
1.
Metals
Found on the
LEFT side of
the PT
2. Nonmetals
Located on the
RIGHT side of
the PT
- Good conductors of heat & electricity
- High melting points most solids at room temperature
- High luster (shiny)
- Ductile (can be drawn into thin wire)
- Malleable (bends without breaking)
- High densities
- Reacts with acids
- Brittle (easy to break)
- No luster (dull)
- Insulators nonconductors
- Neither ductile nor malleable
- Nonreactive with acids
3. Metalloids - Properties of both metals & nonmetals
(Semimetals)
Compounds
 Made up of 2 or more different elements combined in
a fixed position.
 Can be separated through chemical means.
 Represented by formulas.
 Electrolysis allows chemists to distinguish between
elements in compounds.
 Examples: CO2 and H2O
Elements Vs. Compounds
Mixtures: A combination of 2 or more
substances that are not chem. combined.
 Heterogeneous Mixture: Composed of dissimilar components; Can
see the parts
 A.K.A. Mechanical Mixture
 Ex. Cookie, salad, asphalt
 Homogenous Mixture: Uniform structure or composition
throughout
 A.K.A. Solution
 Ex. Lemonade, steel, air
 Alloy: A solid homogeneous mixture (14 caret gold, steel, pewter)
Examples of Alloys
Brass is an alloy of copper
and zinc.
Steel is an alloy of carbon
and iron.
Bronze is an alloy of
copper and tin.
Microscopic look at mixtures
Counting Atoms in Compound
 Step 1: List all elements present
 Step 2: Identify the coefficient
 Step 3: Count the number of atoms of each
element in the compound.
 Step 4: Multiply the coefficient by the subscript
 Step 5: Add up all the atoms
Counting Atoms
 Na2SO4
 Ca(OH)2
 3 Fe2(SO3)3
Separating Heterogeneous Mixtures
 Filtration: Pour liquid through paper and collect residue
(solid)
Separation of Homogeneous Mixtures
 Distillation: Separation based on a difference in boiling
points
Another Look at Distillation
 Distillation Demo
 A Closer Look at Distillation
Separation of Homogeneous Mixtures
 Crystallization: Evaporate liquid and solid will recrystallize
Separation of Homogeneous Mixtures
 Chromatography: Separation of pigments of dye
Percent Concentration of Solutions
 A measure of the amount of solute in a solution.
 % Concentration = mass of solute x 100
mass of solution
 Note: Solute + Solvent = Solution
Solution
 Definition: a homogeneous mixture of 2 or
more substances in a single physical state
 Parts: solute and solvent (usually water)
 Types:
 Physical states: solid (alloys), liquid, gas
 Miscible vs. Immiscible
 Miscible: Liquids that dissolve freely in one another in
any proportion
 Immiscible :Liquid solutes and solvents that are not
soluble
 Saturated, Unsaturated and Supersaturated
 Dilute vs. Concentrated
 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)
Supersaturated Solution of Sodium Thiosulfate
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
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
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 Gases
 Gases are less
soluble at high
temperatures than
at low temperatures
 Increasing
temperature,
decreases solubility.
 Increasing pressure,
increases 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 and Change
 Energy is the capacity to do work.
 All physical and chemical changes require energy.
 Endothermic - describes a process in which heat is
absorbed from the environment.
 Exothermic – describes a process in which heat is
released into the environment.
Law of Conservation of Energy
 Energy is neither created, nor destroyed. It just changes
forms.
Types of Energy
Potential energy – stored energy
Kinetic energy – energy of motion
Heat Transfer
 Transfer of heat may not affect temperature.
 During a phase change, the temperature will
remain constant until all of the substance has
changed state.
 The temperature will increase when a substance is
a solid, liquid, or gas.
Kinetic Theory of Matter
 Gases posses the greatest amount of kinetic energy.
 Two factors that determine the state of matter of a
substance: speed of particles and distance
 There are two factors contribute to the attraction between
the particles.
Kinetic Theory of Matter
 Substances change phases when they
overcome these attractions.
 The overall kinetic energy will not change
until the entire substance has completely
changed.
 Comparison of the three states of matter
Kinetic-Molecular Theory and Gases
1. Gases are small particles that have mass. These particles
are usually molecules, except for the Noble Gases.
2. The particles in gases are separated
by relatively large distances.
3. The particles in gases
are in constant rapid
motion (random).
4. Gases exert pressure
because their particles
frequently collide with
the walls of their
container and each
other.
5. Collisions of gas particles are perfectly elastic.
Inelastic Collision
Elastic Collision
Gas particles do not slow down when hitting each other or
the walls of their container.
6. 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
7. Gas particles exert no force on one another. Attractive
forces are so weak between particles they are assumed to
be zero.
Boyle’s Law
Pressure - Volume Relationship.
The pressure & volume of a sample of gas at constant
temperature are inversely proportional to each other.
Indirect
P1V1 = P2V2
Boyle’s Law
Boyle’s Law Problem
A gas has a volume of 300. mL under a pressure of 740.
mm of mercury. If the temperature remains constant,
calculate the volume when under a pressure of 750.
mm Hg.
P1V1 = P2V2
 740. mm  300. mL =  750. mm   V2 
 740. mm   300. mL
750. mm
=
 750. mm   V2 
750. mm
296 mL = V2
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’ 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
 The pressure of a fixed volume of gas is directly
proportional to its absolute temperature. Law
assumes n is constant.
P1 = P2
T1
Direct
T2
*Temperatures must be in Kelvin!
K = °C + 273
Gay-Lussac’s Law
Gay-Lussac’s Law
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” This law can be used to
determine how changing two
variables at a time affects a third
variable.
P1V1 P2V2
=
T1
T2
A gas occupies 72.0 mL at 25 °C and 198 kPa. Convert
these to standard conditions. What is the new
PV
PV
=
volume?
T
T
P1 = 198 kPa
P2 = 101.325 kPa
V1 = 72.0 mL
V2 = ?
T1 = 298 K
T2 = 273 K
1
1
1
2
2
2
198 kPa  72.0 mL  = 101.325 kPa  V2
298 K
129 mL = V2
273 K
Dalton’s Law of Partial Pressures
 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 Pressures
 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