Matter and Energy

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Matter and Energy

• Unit I

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I. Chemistry

• Definition the branch of science that deals
  with the study of matter, energy, and their
  interaction.

a) Matter defined as anything that has mass and
occupies volume.

• Mass is the measurement of the amount of
  matter in an object OR is the quantity of matter
  an object has.

• The fundamental unit for mass is the gram (g).

• The instrument used in chemistry to measure the
  mass of an object is the triple beam balance or
  the digital scale.

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Figure 1 Triple Beam Balance From
http//www.kirkwood.k12.mo.us/parent_student/khs/B
artinJ/sci20skills20book/using_a_3beam_bal.jpg
4

• Volume defined as the amount of space an object
  occupies.

• The fundamental unit for volume is the liter
  (l). Other units used to measure the volume of
  substances include the cubic centimeter (cm3) or
  the (cc).

• The instrument used in chemistry to measure the
  volume of a substance is the graduated cylinder.

Fig 2 Graduated Cylinder showing the
meniscus From http//www.electrickiva.com/quiz/e
xam03/meniscus.gif
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b) The States of Matter

• Gases, liquids and solids are all made up of
  microscopic particles, but the behaviors of these
  particles differ in the three phases.

1. SOLID state of matter that has BOTH A
DEFINITE VOLUME and a DEFINITE SHAPE. 
Fig 3 Arrangement of Particles in a Solid
From http//www.ul.ie/walshem/fyp/solid.gif
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 2. LIQUID state of matter that has DEFINITE
VOLUME but NO DEFINITE SHAPE.  A key property of
a liquid is that they FLOW and can be POURED.

Fig. 4 Arrangement of Particles in a Liquid
From http//www.ul.ie/walshem/fyp/states20of
20matter.htmliquid
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3. GAS state of matter that has NO DEFINITE
VOLUME and NO DEFINITE SHAPE. 

• A Gas ALWAYS TAKES BOTH THE VOLUME AND THE SHAPE
  OF ANY CONTAINER INTO WHICH IT IS PLACED.  If a
  gas is NOT in a container, it will spread out as
  far as it can.

Fig. 5 Arrangement of Particles in a Gas From
http//www.ul.ie/walshem/fyp/gas.gif
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Ex 1
Under the same conditions of temperature and
pressure, a liquid differs from a gas because the
particles of the liquid a) are in constant
straight-line motion b) take the shape of the
container they occupy c) have no regular
arrangement d) have stronger forces of
attraction between them
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II. Energy
a) Definition The capacity for doing work.

• b) Energy has several forms including
• potential energy - the energy stored in the
  chemical bonds that exist between particles of
  matter.
• kinetic energy the energy of motion.
  Temperature is a direct measure of the average
  kinetic energy of particles.

• c) Law of Conservation of Mass and Energy
• This scientific law states that neither mass nor
  energy can be created or destroyed. They can
  only be converted from one form into another.

• Units Joule (J) or Kilojoules (KJ), calorie
  (c) or Kilocalorie (KJ).

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d) Types of Energy

• Light Energy the energy associated with light
  waves and other forms of electromagnetic
  radiation.

• Electrical Energy the energy associated with
  electrical current (flow of electrons).

• Chemical energy the energy associated with
  chemical changes (the breaking and reformation of
  chemical bonds).

• Heat Energy the energy associated with the
  temperature of substances. Heat energy is the
  least useful form of energy.

• Mechanical Energy the energy associated with
  doing work.

• Atomic and Nuclear Energy the energy associated
  with changes in the mass of atoms and the energy
  that binds atoms together.

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Ex 2 As ice cools from 273 K to 263 K, the
average kinetic energy of its molecules will a)
decrease b) increase c) remain the same
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e) Energy and Chemical Change

• All chemical reactions require energy to occur.

• Chemical bonds are forces of attraction that
  hold atoms, elements, compounds, and molecules
  together.

• Chemical bonds store energy.

• For a reaction to occur, a chemical must absorb
  enough energy to break the bonds that hold its
  atoms together.

• The amount of energy that must be absorbed by a
  chemical to begin a reaction is called the
  Activation Energy (Ea).

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III. Free Energy and Spontaneity
a) Spontaneous Reactions chemical reactions
that occur without the addition of an outside
source of energy.
b) Non-spontaneous Reactions chemical
reactions that require an external energy source
to occur.
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• NOTE
• Enthalpy (H) indicates whether a reaction is
  exothermic or endothermic.

• If DH (), the reaction is endothermic and
  energy is absorbed.

• If DH (-), the reaction is exothermic and energy
  is released.

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• Types of Reactions
• Exothermic Reactions
• Energy releasing processes, ones that "generate"
  energy, are termed exothermic reactions.

Figure 6
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Figure 7 Exothermic Reaction
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• b) Endothermic Reactions
• Reactions that require energy to initiate the
  reaction are known as endothermic reactions.

Figure 8 Endothermic Reaction
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• Note
• All natural processes tend to proceed in such a
  direction that the disorder or randomness of the
  universe increases.

• Endergonic Endothermic
• Exergonic Exothermic

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V. Measuring Energy
a) Temperature

• Temperature is a measure of the average kinetic
  energy of matter.

• The greater the average kinetic energy, the
  greater the velocity of the particles of matter,
  the greater the temperature (and vice versa).

• The instrument used to measure the temperature
  (average kinetic energy) of matter is the
  thermometer.

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Do Now Label the following diagram and indicate
what each arrow stands for.
D
B
E
A
F
C
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1. Thermometer

• A thin glass, capillary tube that contains a
  fluid (mercury (Hg)) that when heated will
  expand.
• When the fluid expands, it rises. This
  correlates to an increase in temperature.

2. Temperature Scales

• There are several scales that are used to
  represent the temperature of substances. They
  include Fahrenheit, Celsius, and Kelvin.

• Scientists most frequently use the Celsius and
  Kelvin scales.

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212
100
373
Boiling Point
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0
273
Freezing Point
0
Absolute Zero
Fahrenheit
Celsius
Kelvin
Figure 9 Temperature Scales and the
Boiling/Freezing Points of Water at Standard
Conditions
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Conversion Formulas C Celsius K Kelvin 1.
C K 273 2. K C 273
Ex 3 What Kelvin temperature is equal to
25C? a) 248 K b) 298 K c) 100 K d) 200 K
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b) Heat Energy

• The device that is used to measure the amount of
  heat energy that a substance contains is called a
  calorimeter.

Figure 10 Calorimeter
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• Substances store energy in their chemical bonds.
  

• When the chemical bonds are broken, heat energy
  is released.

• The heat energy that is released by the
  substance should theoretically be equally to the
  chemical energy stored in the chemical bonds of
  that substance.

• A calorimeter indirectly determines the amount
  of heat energy stored by a substance by measuring
  the change in temperature for a known quantity of
  water which immerses that substance.

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1. Determination of the Heat Energy Transfer

• To determine the amount of heat energy
  gained/lost by a substance, several variables
  must be known. These include

Q heat energy lost/gained (Joules) m mass
(grams) c specific heat capacity (J/g. C o) D
T change in Temp
C water 4.2 J/g. C o
Reference Table B
Q mc D T
Reference Table T
Note Specific heat capacity indicates the ease
at which a substance absorbs or releases heat
energy. It is a constant value for any given
substance.
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Figure 11 Calorimeter
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Ex 4 What is the total number of joules of
heat energy absorbed by 15 grams of water when it
is heated from 30C to 40C? a) 10 b) 63 c)
150 d) 630
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2. Determination of the Amount of Energy to
Required for a Phase Change
a) Heat of Fusion the amount of energy required
to convert a solid to a liquid at its melting
point and standard pressure.
Q heat energy lost/gained (Joules) m mass
(grams) Hf Heat of Fusion (for water -
reference table B)
Q mHf
334 J/g
Q mHv
Q heat energy lost/gained (Joules) m mass
(grams) Hv Heat of Vaporization (for water -
reference table B)
2260 J/g
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Do Now Calculate the following thermochemistry
problem.
1. Find the initial temperature of a 120g sample
of water that has reached a final temperature of
90ºC by absorbing 2520J of energy. NOTE C
water 4.2 J/g. C o

• A 10 g sample of water at -20º C is constantly
  heated till it reaches a final temperature of
  110º C. Find the amount of heat energy needed
  for this change to occur given the following
  constants
• C ice 1.96J/g. C
• C water 4.2 J/g. C
• Hf 334 J/g
• Hv 2260 J/g

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VI. The Classification of Matter

• a) Substances Elements and Compounds
• 1. ELEMENTS ARE PURE (homogeneous) SUBSTANCES
  THAT CANNOT BE BROKEN DOWN (decomposed)
  CHEMICALLY INTO SIMPLER KINDS OF MATTER.

• More than 100 elements have been identified,
  though Fewer than 30 are Important in Living
  Things. 

• All of the Elements are arranged on a Chart
  known as THE PERIODIC TABLE. 
• Among the information provided in The Periodic
  Table are the ATOMIC NUMBER, THE CHEMICAL SYMBOL,
  AND THE ATOMIC MASS FOR EACH ELEMENT.

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• More than 90 Percent of the Mass of living
  things is composed of JUST FOUR ELEMENTS 
  OXYGEN, O, CARBON, C, HYDROGEN, H, AND NITROGEN,
  N.

• Each Element has different Chemical Symbol which
  consist of One or Two Letters.

Figure 12 Atomic Mass, Number, and Chemical
Symbol
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2. Compounds Substances that consist of two or
more elements that are combined chemically by
bonds.

• The elements in a compound can only be separated
  chemically by breaking the bonds that hold them
  together.

• Compounds have fixed ratios of their components.
  Water will always have 2 hydrogen atoms for every
  one oxygen atom.

• Compounds are homogeneous one cannot
  distinguish between the components of the
  compound. If given a sample of water, you could
  not determine what is hydrogen and what is
  oxygen.

Examples of Compounds
water
carbon dioxide
glucose
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• Two or more substances that are combined
  physically the components can be easily
  separated.

b) Mixtures

• In homogeneous mixtures, the substances are
  completely mixed. This means that you cannot see
  the individual components. The mixture appears to
  be only one substance.

• In heterogeneous mixtures, the substances are
  not completely mixed. This means that you can see
  the individual components. The mixture appears to
  be only two different substances in the same
  container.

Figure 14 Heterogeneous Mixture SOIL
Figure 13 Homogeneous Mixture Salt Water
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Ex 5 A compound differs from a mixture in that
a compound always has a a) homogeneous
composition b) maximum of two components c)
minimum of three components d) heterogeneous
composition
Ex 6 Which substance cannot be decomposed into
simpler substances? a) ammonia (NH3) b) aluminum
(Al) c) methane (CH4) d) methanol (CH3OH)
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VII. Heating/Cooling Curves

• Heating/Cooling curves shows the change in
  kinetic and potential energies of substances. In
  addition, the curves indicates the points of
  phase change.

a) Kinetic Energy the energy associated with
the velocity of the particles of matter.
Temperature is a direct measure of the average
kinetic energy of matter. As temperature
increases, so does the average kinetic energy of
matter (and vice versa).

• Note During phase changes, the KE of matter
  remains the same.

b) Potential Energy the stored energy found
within substances. The potential energy of a
substance increases as it is converted to a phase
of matter having greater entropy (and vice
versa).

• Note When KE changes, PE remains the same.

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Figure 15 The Heating Curve of Water
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Explanation of Figure 15
Point State Kinetic Energy Potential
Energy A Solid Increases RTS B (s) to (l)
melting RTS Increases C Liquid Increases RTS
D (l ) to (g) boiling RTS Increases E Gas
Increases RTS
RTS remains the same
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• c) Phase Changes
• Substances must either gain or lose energy for a
  change of phase to occur.

Increase in Enthalpy (and Entropy) Process Solid
to liquid Melting Liquid to gas Boiling
(vaporization) Solid to gas Sublimation
Decrease in Enthalpy (and Entropy) Process Gas to
Liquid Condensation Liquid to
solid Freezing Gas to solid Deposition
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d) Heats of Fusion and Vaporization

• Heat of Fusion the amount of energy needed to be
  gained by a substance to convert it from a solid
  to a liquid.
• This is an intensive property.
• Example Water - Hf 334 J/g

• Heat of Vaporization the amount of energy
  needed to be gained by a substance to convert it
  from a liquid to a gas.
• This is an intensive property.
• Example Water - Hv 2260 J/g

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e) Heats of Fusion/Vaporization Calculations
Q mHf Q mHv
Q heat energy in joules m mass in grams Hf
Heat of Fusion Hv Heat of Vaporization
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Ex 6 How many joules of heat are absorbed when
70.0 grams of water is completely vaporized at
its boiling point? a) 23, 352 J b) 7, 000 J c)
15, 813 J d) 158, 130 J
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VIII. The Unique Properties of Gases
a) Gases are greatly influenced by changes in
temperature and pressure.
b) The behavior of gases is due in part to the
fact that the atoms/molecules that they are
comprised of are greatly dispersed.
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c) The Kinetic-Molecular Theory of Gases (Ideal
Gases)
1. Gases are made from molecules that are in
constant random motion.
2. Collisions between gas particles are
completely elastic.
3. The volume of individual gas molecules is
insignificant as compare to the overall volume of
space that the gas occupies.
4. Individual gas molecules do not have
attractive forces for each other.
5. The average kinetic energy of a gas is
directly proportional to its temperature.
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d) Exceptions to Ideal Gases
1. As opposed to an ideal gas, real gas molecules
do have small but significant volumes as well as
attractive forces.
2. These deviations become apparent under the
conditions of low temperature and high pressure.
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e) The Gas Laws

• Boyles Law an indirect relationship states
  that at a constant temperature the volume of a
  gas decreases with increasing pressure.
• P1V1 P2V2

KEY P1 initial pressure P2 final
pressure V1 initial volume V2 final volume
Figure 16 Boyles Law From http//chemmovies.u
nl.edu/chemistry/smallscale/SSGifs/26Fig2.gif
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2. Charles Law a direct relationship states
that at a constant pressure, the volume of a gas
will increase with a corresponding increase in
the temperature (Kelvin). V1T2 V2T1
KEY V1 initial volume T2 final temperature V2
final volume T1 initial temperture
Figure 17 Charles Law From http//wine1.sb.fsu
.edu/chm1045/notes/Gases/GasLaw/charles.gif
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• Combined Gas Law
• (P1) (V1) (P2) (V2)
• T1 T2

NOTE Standard Temperature and Pressure 0 C
or 273 K 101.3 kPa or 1 atm
http//www.chem.iastate.edu/group/Greenbowe/sectio
ns/projectfolder/animationsindex.htm
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Vapor Pressure

• Vapor pressure is the pressure of a vapor in
  equilibrium with its non-vapor phases. All solids
  and liquids have a tendency to evaporate to a
  gaseous form, and all gases have a tendency to
  condense back. At any given temperature, for a
  particular substance, there is a partial pressure
  at which the gas of that substance is in dynamic
  equilibrium with its liquid or solid forms. This
  is the vapor pressure of that substance at that
  temperature.

• Equilibrium vapor pressure is an indication of
  a liquid's evaporation rate. It relates to the
  tendency of molecules and atoms to escape from a
  liquid or a solid. A substance with a high vapor
  pressure at normal temperatures is often referred
  to as volatile. The higher the vapor pressure of
  a material at a given temperature, the lower the
  boiling point.
• The boiling point of a liquid is the temperature
  where the vapor pressure equals the ambient
  atmospheric pressure. At the boiling temperature,
  the vapor pressure becomes sufficient to overcome
  atmospheric pressure and lift the liquid to form
  bubbles inside the bulk of the substance.

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VIII. Avogadros Hypothesis

• Under the same conditions of temperature and
  pressure, equal volumes of all gases contain the
  same number of particles.

• For example, 1 liter of hydrogen gas will
  contain the same number of particles as 1 liter
  of oxygen gas.