CHEMISTRY REGENTS REVIEW
SPRING 2012
PHYSICAL BEHAVIOR OF MATTER
WEEK # 4
REGENTS TEXTBOOK:
HONORS TEXTBOOK:
BARRONS REV. BOOK:
CHAPTER 2, 13, & 14
CHAPTER 3, 13, & 14
CHAPTER 1, 6
CHEMISTRY REGENTS
WED.JUN 20, 2012
12:15 PM
Section I: Matter
I.
II.
Chemistry - the study of the composition, structure, and properties of matter, the
changes matter undergoes and the energy accompanying these changes.
A. Branches of Chemistry
1.
Organic Chemistry – study of carbon – hydrogen containing
compounds
2.
Inorganic Chemistry – study of noncarbon compounds
3.
Physical Chemistry – study of the properties and changes of matter
and their relation to energy
4.
Analytical Chemistry – identification of the components and
composition of materials
5.
Biochemistry – study of substances and processes occurring in
living things
6.
Theoretical Chemistry – use of mathematics and computers to
understand the principles behind observed chemical behavior and
to design and predict the properties of new compounds.
Matter - anything that has mass and takes up space (volume)
A.
B.
Pure Substance - homogeneous matter having identical properties and
composition.
1.
Element - composed of only atoms of the same atomic number
a.
cannot be decomposed by chemical means
2.
Compound - two or more different elements chemically combined
in a definite ratio by weight (using atomic mass)
a.
chemical and physical properties are different than the
elements that make up the compound
b.
can be made from simpler substances
c.
can be decomposed into its separate elements
1.
Binary Compound - contains 2 types of elements
2.
Ternary Compound - contains 3 types of elements
Mixtures - combinations of varying amounts of 2 or more distinct
substances
1.
Homogeneous Mixture - uniform intermixture of particles when
one substance dissolves in another (solution)
a.
gas in gas - air
b.
solid in liquid - salt in water
c.
solid in solid - an allow such as brass (copper & zinc)
d.
liquid in liquid - alcohol in water
2.
Heterogeneous Mixture - have uniformly dispersed ingredients.
Two or more phases can be seen.
Examples:
a.
concrete
b.
sand and water
c.
oil and water
III.
IV.
Properties - a definite set of characteristics by which a substance can be identified.
1.
can be observed by examining the substance
2.
Determined by the manner in which it behaves when in contact
with other substances or sources of energy.
B.
Extensive Properties - depend upon the quantity of a substance present
1.
volume, weight, mass
C.
Intensive Properties - do not depend on size
1.
melting point, boiling point, density,
D. Physical Properties - characteristics which can be observed without
producing new substances
1.
Physical change - no new substance is formed but changes of
phase may occur.
a.
Ex. Grinding, freezing, boiling
2.
Phases of Matter
a.
Solid – definite volume and definite shape
b.
Liquid – definite volume but and indefinite shape (takes
shape of container)
c.
Gas – no definite shape or volume (completely fills any
container)
d.
Plasma – high temperature physical state of matter in which
atoms lose their electrons
E.
Chemical Properties - describe how a substance interacts (or not) with
other substances
1.
Chemical change (reaction) - results in the production of 1 or more
new substances
a.
rusting of iron, burning paper
F.
Changes in Energy - occur in both physical and chemical changes
1.
can be either absorbed or released
G. Law of Conservation of Matter or Mass (Lavoisier) - matter cannot be
created or destroyed by a chemical change
Physical Phases (states) of Matter
A. Solid - the substance is relatively rigid and has a definite volume and
shape
1.
must have a "precise" temperature at which it melts
2.
particles are in a fixed position with little space between them
a.
these particles (atoms or molecules) are vibrating)
B.
Liquid - a substance has a definite volume, but its shape changes by
flowing
1.
Particles are considered to vibrate and rotate
a.
Ex. glass becomes softer as it is heated & flows, but does
not change phase- glass is considered a liquid with a high
viscosity
b.
Viscosity - resistance of a liquid to flow
C.
D.
Gas - a substance has no definite volume nor shape and has little response
to gravity
1.
Particles are considered to vibrate, rotate and translate
a.
Translating – particles break intermolecular bonds allowing
for random motion
Plasma – a high temperature state in which atoms lose their electrons
a.
Transition metals – Group 11 is +1, +2 (except for silver
(Ag) which is +1 only
Section II: Liquids & Solids
I.
Liquid – form of matter that has a definite volume and takes the shape of its
container.
A. Application of KMT
1.
liquid particles are closer together than gas and have a lower
kinetic energy
2.
attractive forces between particles is greater than that of a gas
3.
IMF’s determine the temperature at which condensation or
liquifactiion occurs dependent upon the H-bonding, dipole dipole
attractions or VanderWaal’s (London Dispersion) forces acting
upon the substance
B.
Pressure changes have little or no effect on liquids
C.
Cohesion – ability of a liquid to be attracted to other molecules of itself
D. Adhesion – ability of a liquid to be attracted to molecules of another
substance
1.
II.
Capillary action – ability of a liquid to move up a tube of a solid
substance due to IMF’s
Boiling Point and Melting (freezing) Point
A.
Boiling Point - temperature at which a substance changes phase between
liquid and gas at a particular pressure
1.
ex. water can boil at various temperatures depending upon the
pressure
(>1Atm = higher boiling point,
<1Atm = lower boiling point)
2.
Normal Boiling Point = vaporization at 1Atm (ex. water is 100C
at 1Atm)
B.
C.
D.
E.
F.
G.
III.
Melting (freezing) Point - temperature at which a substance makes a phase
change between the liquid and solid phases
1.
Normal Melting Point = phase change at 1Atm (ex. water is 0Cat
1Atm)
Evaporation - process in which a liquid changes to a gas at the surface of
the liquid
1.
happens at any temperature in liquid phase (forms a vapor)
a.
Increasing rate of evaporation
1.
raise temperature of liquid
2.
increase surface area
3.
air currents over surface to carry away molecules
2.
tends to lower the temperature of the liquid
Boiling - liquid changes to a gas throughout the liquid
1.
temperature at which the vapor pressure is equal to the pressure on
the surface of a liquid
a.
low pressure - low boiling temperature
b.
high pressure - high boiling temperature
2.
Different substances have different boiling points (due to different
vapor pressures)
Vapor Pressure - the vapor in equilibrium with its liquid (boiling points at
different pressures)
1.
Reducing the pressure will allow a liquid to boil at a lower
temperature
**Table G in the reference tables
2.
Substances have different vapor pressures
a.
higher vapor pressure, lower boiling point
1.
more kinetic energy - molecules escape easier
Condensation - process in which a vapor or gas changes to a liquid
1.
lower energy molecules return to the liquid stage
2.
raises the temperature of the vapor
Distillation - process of evaporating off a liquid and recollecting it by
cooling in another container - used to collect pure samples of substances
Solids
A. Crystals - all true solids form characteristic geometric figures in which the
atoms or molecules are arranged in regular repeating patterns
1.
Rate of evaporation (cooling) determines size of crystal formation
a.
slow - atoms attach themselves to previously formed
crystals
b.
fast - atoms form their own centers of crystallization
2.
Types: Ionic crystals, Covalent Network Crystals, Metallic
Crystals, Covalent Molecular Crystals (See Bonding)
B.
C.
D.
E.
Water of Hydration (Hydrates)
1.
Crystals made up of solid substance combined chemically with
water in a definite ratio. (hydrate)
a.
Example CuSO4 5H2O (1 unit of copper sulfate is
chemically united with 5 water molecules)
1.
Anhydrous - removal of water from hydrates
2.
Efflorescence - spontaneous loss of water of
hydration from a substance at room temperature
(ex. Na2CO3  10H2O)
Heat of fusion - same as the Heat of Crystallization
1.
quantity of heat given up at a constant temperature, when one gram
of liquid is changed to a solid.
Sublimation
1.
solids with a fairly high vapor pressure and low intermolecular
attractions will sublimate
a.
solid CO2
b.
iodine
c.
naphthalene (moth balls)
Deposition – process by which a gas changes directly to a solid without
passing through the liquid phase
Section III: The Gas Laws and Kinetic
Theory
I.
Kinetic - Molecular Theory of Gases (particles of matter are always in motion)
A. Gases are composed of separate, tiny particles called molecules.
1.
Gas molecules possess vibrational, rotational and translational
movement
2.
molecules are very far apart
3.
molecules are very small
4.
Gas consists mostly of empty space (volume is determined by the
container)
a.
gases mix easily
b.
gases can be easily compressed
5.
Motion causes collisions with other molecules and walls of
container
a.
collisions result in creating pressure
B.
Ideal gas – an imaginary gas that perfectly fits all the assumptions of a
kinetic-molecular theory
1.
C.
D.
E.
Gases consist of large numbers of tiny particles that are far apart
relative to their size.
2.
Collisions between gas particles and between particles and
container walls are elastic collisions
a.
Elastic collision – one in which there is no net loss of
kinetic energy
3.
Gas molecules are in constant, rapid, straight-line motion (kinetic
energy)
4.
No forces of attraction or repulsion between gas particles
5.
Average kinetic energy (temperature) of the gas particles depends
on the temperature of the gas
a.
KE = ½ mv2
1.
m= mass in Kg
2.
v= velocity in m/s
Effects on real gas molecules
1.
Most gases obey the gas laws up to a point
2.
High pressure or low temperatures cause deviations from this
"ideal"
a.
High pressure causes a higher than expected volume
(constant temperature)
1.
intermolecular forces will decrease the volume at
first
2.
with tremendous pressure, volume becomes a factor
and molecules push out causing a higher than
expected volume
b.
High pressure and lower temperature causes lower than
expected volume
c.
Intermolecular forces and closeness of the molecules are
the cause of the deviations
The average kinetic energy of the molecules is directly proportional to the
Kelvin temperature of the gas
1.
Molecules have different velocities (different kinetic energy)
Pressure - force per unit area (Pressure = force)
a.
Atmospheric Pressure (air) = 1kg / cm2 at sea level
b.
Measuring Air Pressure - mercury barometer
c.
Normal atmospheric pressure (measured at sea level) = 1
atmosphere = 760mm Hg (standard pressure)
d.
Pressure of 1mm of Hg = 1 torr therefore 1atm. = 760 torr
2.
Unit of pressure in the metric system is the pascal
a.
101.3 kpascal (kPa) = 1Atm
F.
II.
III.
IV.
V.
VI.
Measuring Gas Pressure
1.
Manometer - U tube device filled with mercury which changes in
comparison to normal air pressure
2.
Molecular collisions are completely elastic
3.
No kinetic energy is changed into other energy forms but may
transfer between molecules of gas.
a.
Total kinetic energy remains the same with constant
temperature and volume.
1.
Increase temperature - causes increase in pressure
2.
Decrease temperature - causes decrease in pressure
3.
Increase volume - causes decreases in pressure
4.
Decrease volume - causes increase in pressure
Standard Temperature and Pressure
A. Conditions of temperature and pressure that have been agreed upon for the
purpose of determining the mass of a gas
1.
Temperature = 0C (273K)
2.
Pressure = 760mm Hg (760 torr) (1 atm)
B.
Standard conditions for liquids and solids are 25C (298K) and 1Atm
C.
STP allows us to use volume to determine the mass of a gas
1.
Mass will vary directly with the volume (density)
Boyle’s Law - the volume of a sample of gas is inversely proportional to the
pressure if the temperature is kept constant.
A. PV = k (constant)
1.
a sample of gas with fixed mass and constant temperature will
have a constant value regardless of the pressure or volume changes
a.
P1V1 = P2V2
Charle’s Law - at constant pressure, the volume of a gas is directly proportional to
its Kelvin temperature.
1.
V/T = k (k is a proportionality constant produced when V is
divided by T)
2.
The constant changes if
a.
the mass changes
b.
the kind of gas is different
c.
the pressure changes
3.
V1 / T1 = V2/ T2
4.
Also applied is P1 / T1 = P2 / T2
Gay Lussac’s Law – the pressure of a fixed mass of gas at constant volume varies
directly with the Kelvin temperature
A. P/T = k
Combined Gas Laws - allows us to determine the effect of pressure, volume and
temperature on a gas.
A. P1V1 / T1 = P2V2 / T2
1.
If the pressure and volume of a given gas are multiplied, and the
resulting product is divided by the temperature, the result will be a
constant for all combinations of the factors.
VII.
Dalton’s Law of Partial Pressures - In a mixture of gases, the total pressure of the
mixture is equal to the sum of the pressures that each gas would exert by itself in
the same volume
A. Partial pressure of each gas is determined by its molecular ratio (mole
ratio)
1.
Total pressure = Pressure of gas A + Pressure of Gas B + Pressure
of Gas C
2.
= 1.0 atm + 2.0 atm + 3.0 atm
3.
= 6.0 atm
B.
Gases
1.
Collection of Gases - done over water
a.
Resulting gas mixture - combination of gas and water vapor
b.
must subtract partial pressure of water vapor at room
temperature from total temperature to determine pressure of
gas
Gas Stoichiomentry
A.
IX.
Avogadro’s Hypothesis - equal volumes of different gases, at the same
temperature and pressure, contain the same number of molecules
B.
Number of molecules = 6.02 X 1023 at STP (mole)
1.
A mole of particles of any gas occupies a volume of 22.4 L at STP
2.
The number of particles is equal, however the mass of the sample
and the size of the molecules varies with substance.
C.
Mole - the unit used to measure the number of particles (atoms or
molecules) of any kind
1.
based on atomic mass or molecular weight
D. Ideal Gas Law - takes into account the size of any sample size of a gas.
Graham’s Law of Diffusion - under constant conditions of temperature and
pressure, gases diffuse at a rate inversely proportional to the square roots of their
densities. The heavier the gas, the slower its rate of diffusion.
Matter: A Quick Review
1. Matter is classified as a pure substance or a mixture of substances.
 A substance has fixed composition and uniform properties throughout the
sample. Element and compounds are substances.
 A substance MUST be homogeneous.
2. A mixture is composed of two or more different substances that may be physically
separated.
 A mixture may be homogeneous (uniform – a solution), or heterogeneous
(uneven).
 Substances in a mixture retain their original properties.
 Substances in a mixture may be separated by their size, polarity, density, boiling
and freezing points, and solubility (among others).
 Filtration and distillation are examples of processes used to separate mixtures.
3. An element is a substance composed of atoms with the same atomic number. They
cannot be broken down by chemical change.
4. A compound is two or more elements bonded together. It can only be broken down
by chemical changes.
 Substances that form a compound gain new properties.
 The ratio of substances in a compound is constant (e.g. water has a fixed ratio
2:1 ratio of hydrogen to oxygen).
5. A physical change is one that results in the rearrangement of existing particles in a
substance (ex: freezing, boiling). A chemical change results in the formation of different
substances with different properties.
 Chemical and physical changes may be endothermic or exothermic.
6. The three phases of matter are solid, liquid and gas. Each has its own properties.
 Solids have a constant volume and shape. Particles are held in a rigid, crystalline
structure.
 Liquids have a constant volume but a changing shape. Particles are mobile but
still held together by strong attraction.
 Gases have no set volume or shape. They will completely fill any closed
contained. Particles have largely broken free of the forces holding them
together.
7. A heating curve (or cooling curve) traces the changes in temperature of a substance
as it changes from solid to liquid to gas (or gas to liquid to solid).
 When the substance undergoes a phase change, there is no change in
temperature. The line “flattens” until the phase change is complete.
 When a phase change is occurring, the potential energy of the substance
changes while kinetic energy remains the same.
 As temperature increases, kinetic energy increases.
8. Heat of fusion (Hf) is the energy needed to convert one gram of a substance from
solid to liquid.
9. Heat of vaporization (Hv) is the energy needed to convert one gram of a substance
from liquid to gas.
10. Specific heat (C) is the energy required to raise one gram of a substance 1 degree
(Celcius or Kelvin).
 The specific heat of liquid water is 1 cal/g*J or 4.2 J/g*K.
11. The combined gas law states the relationship between pressure, temperature and
volume in a sample of gas.
 Increasing pressure causes a decrease in volume (inverse relationship).
 Increasing temperature causes an increase in volume (direct relationship).
 Increasing temperature causes an increase in pressure.(direct relationship).
12. An ideal gas model is used to explain the behavior of gases. A real gas is most like an
ideal gas when it is at high temperature and low pressure.
 The real gases that most resemble ideal gases are H2 and He
13. The Kinetic Molecular Theory (KMT) for an ideal gas states that all gas particles:
 are in random motion.
 have no forces of attraction between them.
 have a negligible volume compared to the distances between them.
 have collisions that result in the transfer of energy from one particle to another,
with no net loss of energy from the collision.
14. Equal volumes of gases at the same temp and pressure have an equal number of
particles.
Section IV:
PHYSICAL BEHAVIOR OF MATTER: MAJOR POINTS
1.
What is a pure substance? What are the two types of pure substances: How
can they be represented by particle diagrams?
2.
How are the characteristics of solids, liquids, and gases different?
3.
How are the particles arranged differently in solids, liquids, and gases?
4.
What is a mixture: What are the two different types of mixtures? How
would you represent a mixture in a particle diagram?
5.
How do physical properties affect the separation of a mixture? Explain
different separation techniques.
6.
What is heat: What is temperature? How are they related and different:
What are the different temperature scales, and how are they related?
7.
Draw and understand heating and cooling curves. What are the various
phase changes and points?
8.
Which are exothermic processes/ Which are endothermic processes:
9.
10.
How does the kinetic and potential energy change for a substance that is
cooling and eventually solidifies? For a substance that is heating and
eventually vaporizes/boils?
How is a chemical change different from a physical change?
SECTION V: VOCABULARY
BAROMETER___________________________________________________________
BOILING POINTS_______________________________________________________
BOYLE’S LAW__________________________________________________________
CHARLES’S LAW________________________________________________________
COMBINED GAS LAW___________________________________________________
COMPRESSIBILITY_____________________________________________________
DALTON’S LAW OF PARTIAL
PRESSURE_____________________________________________________________
DIFFUSION____________________________________________________________
EVAPORATION_________________________________________________________
GAS____________________________________________________________________
GAS PRESSURE_________________________________________________________
GAY-LUSSAC LAW______________________________________________________
GRAHAM’S LAW OF
DIFFUSION____________________________________________________________
KINETIC ENERGY_______________________________________________________
KINETIC THEORY_______________________________________________________
LIQUID________________________________________________________________
MELTING POINT________________________________________________________
NORMAL MELTING/BOILING POINT______________________________________
PASCAL________________________________________________________________
PHASE_________________________________________________________________
SOLID_________________________________________________________________
SOLUTION_____________________________________________________________
SUBLIMATION__________________________________________________________
VAPOR_________________________________________________________________
VAPORIZATION_________________________________________________________
VAPOR PRESSURE______________________________________________________