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Chemistry is the study of chemicals and their properties, and the reactions they undergo.
Matter has mass and volume:
Matter
Pure Substances
Compound or Element
(2+ elements) (unique #p+)
or
Mixtures
Homogenous or Heterogenous
(uniform)
(non-uniform)
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Physical properties/changes involve changes in state, color, etc. that do not involve a change in
the identity of a chemical
 Chemical properties/changes describe a chemicals potential to react, and the changes in identity.
 Evidence of a chemical reaction: color change, gas evolution, precipitate formation, release of
energy (heat)
 Density is mass per unit of volume (g/ml)
 Mass is the amount of matter; weight is the force of gravity on an object
 SI units (Systeme Internationale); the rest of our units are derived from these basic units
length
l
meter
m
mass
m
kilogram
kg
time
t
second
t
temperature
T
kelvin
K
amount of substance n
mole
mol
electric current
I
ampere
A
luminosity
Iv
candela
cd
Derived Units
force
newton
N
m·kg·s-2
2
pressure, stress
pascal
Pa
N/m
m-1·kg·s-2
energy, work, quantity of heat
joule
J
N·m
m2·kg·s-2
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Conversions: the key to any conversion is to determine first what you have and then what you
need or want. To set up the problem, simply multiply the data you are given by the conversion
factor. The units of the conversion factor are what you need over what you have.
 Metric table
Mega
M
106
1,000,000
3
Kilo
K
10
1,000
0
Base
10
1
Deci
d
10-1
0.1
-2
Centi
c
10
0.01
-3
Milli
m
10
0.001
-6
Micro
10
0.000 001
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Energy is the capacity to do work and comes in 6 forms: chemical, electrical, light, mechanical,
sound, heat
Heat is the exchange of energy from high to low; temperature is the measure of how hot or cold
something is
 Law of Conservation of Energy: energy is conserved
 Heat of Fusion and Heat of Vaporization are at the melting and boiling points of a compound
respectively
 Specific Heat is the amount of heat needed to raise one gram of a compound by 1 degree
Celsius or Kelvin (cp=q/m.T)
 Significant figure rules:
o Non-zero digits are always significant.
o Any zeros between two significant digits are significant.
o A final zero or trailing zeros in the decimal portion ONLY are significant.
 0.0001 has only 1 significant figure
 0.00010 has 2 significant figures
 1.0001 has 5 significant figures
 200 has only 1 significant figure
 200. has 3 significant figures
o When quantities are added or subtracted, the number of decimal places in the answer is
equal to the number of decimal places in the quantity with the smallest number of
decimal places
o In multiplication and division, the result may have no more significant figures than the
factor with the fewest number of significant figures.
 Accuracy is how close data is to the true value; Precision is how close the data in repetitive
trials are
 Endothermic reactions absorb heat to proceed; Exothermic reactions release heat
o Reactant + heat  Product
decrease in temperature of surroundings
o Reactant  Product + heat
increase in temperature of surroundings
 System: the entire group or sample set; Surroundings: the environment outside of the system
 Law of Conservation of Mass: Mass is conserved
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Democritus: Atoms are the smallest fraction of a molecule
Three Laws of Atoms
o Law of Definite Proportions
o Law of Multiple Proportions
o Law of Conservation of Mass
Dalton’s Atomic Theory:
o All matter is composed of atoms that cannot be subdivided, created or destroyed
o Atoms of an element have the same properties
o Atoms of different elements have different properties
o Compounds are formed when different atoms combine in simple, whole numbers
o In a chemical reaction, atoms combine, separate or rearrange, but are not created,
destroyed or changed
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Atoms are made up of subatomic particles called protons, neutrons and electrons
J.J. Thomson set up a crookes tube with a anodic and cathodic ends
o When electricity was applied to the tube, a beam was emitted from the cathodic (-) plate
o He tested the tube further by applying an electrical field to the tube using paddles
o He concluded that the particles in the tube were negatively charged and had mass
 electrons, e, e-, -1.602 x 10-19C; mass = 9.109 x 10-31kg
o J.J. Thomson’s Plum Pudding Model
Robert Milikan calculated the charge of an electron
Ernest Rutherford conducted experiments to test the Thomson model
o He directed alpha particles through a thin gold foil and measured them with a film
o There must be a dense region with positive charges (protons) surrounded by the
electrons; nucleus
o Rutherford concluded that the nucleus had all of the positive charge and most of the
mass, but only a fraction of the volume
o Most of the volume of an atom is empty space
The Atomic Number is the number of protons in the nucleus and is unique for each element
o The Atomic Mass Number is the number of Protons + Neutrons
o Atomic Mass Number = Atomic Number + Number of Neutrons
Different elements can have the same mass number, but different atomic number
You cannot use mass number to identify an element
The structure of the atom can be written using the element symbol, atomic number and mass
number
An Isotope has the same atomic number but different mass number (different number of
neutrons)
o Isotopes can be identified with the symbol, mass and atomic number
o 3/2He or Helium-3
The Rutherford model described electrons as a cloud of negatively charged particles
surrounding the nucleus
If electrons are negatively charged, and the nucleus is positively charged, why don’t the
electrons collide with the nucleus? Opposite charges attract?
Niels Bohr was a quantam physicist who described energy levels of an electron
o The Bohr Model is probably familar as the "planetary model" of the atom, the figure is
used as a symbol for atomic energy
o In the Bohr Model the neutrons and protons occupy a dense central region called the
nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun
Louis deBroglie – responsible for connecting the wave theory of particles to the modern atomic
theory
Schrodinger Wave Theory: wave-particle duality; electrons exhibit both particle and wave like
properties
Plank stated that energy is absorbed and released in discrete units called quanta
An orbital is a region of space that an electron is most likely to occur
o S-orbitals: are spherical in shape and there is only 1; there are on each principal energy
level
o P-orbitals: there are 3 types of p-orbitals starting on the 2nd level
o D-orbitals: there are 5 types starting on the 3rd level
o F-orbitals: there are 7 types starting on the 4th level
o Think of orbitals as different types of rooms, each room can only hold 2 electrons
Floor #
Number of
Maximum Number of
(principal energy
Type of Orbitals
Orbitals
Electrons
level)
1
S
1
2
2
S, P
4
8
3
S, P, D
9
18
4
S, P, D, F
16
32
2
2
n
n types
n orbitals
2n electrons
Electron Hotel
4th Floor
3rd Floor
2nd Floor
1st Floor
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
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F orbital types
D orbital types
P orbital types
S orbital types
14 electrons +
10 electrons +
6 electrons +
2 electrons =
32 electrons total
D orbital types
P orbital types
S orbital types
10 electrons +
6 electrons +
2 electrons =
18 electrons total
P orbital types
S orbital types
6 electrons +
2 electrons =
8 electrons total
2 electrons
S oribital types
Quantum numbers are unique numbers assigned to each individual electron which indicate the
location and the energy of an electron
 The principal quantum number (n = 1, 2, 3, 4 ...) is the principal energy level
Hund’s Rule: electrons entering an orbital-type (sublevel) will half-fill the orbitals in the sublevel
before they fill it completely
Pauli’s Exclusion Principle: no two electrons can exist in the exact same state; with the same
quantum number
Aufbau’ Principle: electrons occupy the lowest possible energy level
An electron configuration is the arrangement of electrons in an atom
 The diagonal rule: fill each orbital type according to the direction of the arrows
6s
6p
6d
5s
5p
5d
5f
4s
4p
4d
4f
3s
3p
3d
2s
2p
1s
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Henry Mosely – periodic law; elements are arranged in the periodic table by their atomic number =
# of protons = # of electrons
The most modern form of the periodic table was introduced by Mendeleev.
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Main group elements:
 Group 1: Alkali Earth Metals
 Group 2: Alkaline Earth Metals
 Group 17: Halogens
 Group 18: Noble Gases
 Groups 3-12: Transition Metals
 Lanthanides and Actinides are the “f” orbital type elements
Metals are on the left; non-metals are on the right; metalloids are along the stairstep
Trends
 Ionization Energy, Ionic Radius and Electronegativity
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Elements are arranged in columns (group or families) and rows (periods)
Families have similar characteristics; they are said to have periodicity
New elements can be predicted based on these trends
Atomic Radius
Ionic Compounds are formed from a metal and non-metal; the difference in electronegativity is
greater than 2.1
 Most ionic compounds are solid at room temperature because of the strength of the bond
 An ionic compound contains repeating units or formula units bound together in a crystal lattice
 Lattice energy is the amount of energy needed to form or break an ionic bond
 Ionic compounds dissociate in water to form ions; these ions conduct electricity
 Electrons are exchanged in order to form a compound
 Ionic compounds are named with the cation (“+”) and then the anion (“-“)
 The anion is the name of the element + “-ide”
 If the anion is a polyatomic ion, simply name the ion
Ions and parent atoms have different properties
Octet rule; rule of 8: all compounds want to be more like a noble gas, they want an electron
configuration like a noble gas (8 valence electrons).
 Therefore, they will lose/gain electrons to be stable
Covalent Compounds are formed from a metal to metal or a non-metal and non-metal; the
difference in electronegativity is less than 2.1
 Polar covalent bonds are between 0.5 and 2.1
 Non-polar covalent bonds are less than 0.5
 Electrons are shared in order to form a compound
 As bond length increases, bond energy decreases
 A triple bond is the shortest and strongest of all the bonds
 Bond energy is the amount of energy needed for form or break a covalent bond
 Covalent compounds are named with the less electronegative atom first followed by the more
electronegative atom. Prefixes are used to indicate ratios of atoms
Lewis dot structures illustrate the configuration of the valence (outer shell) electrons
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The shape of a bond can predict the overall polarity of a compound (VSEPR or Valence Shell
Electron Pair Repulsion)
 If there is a dipole, or an uneven shift of charge (remember that a compound cannot have an
overall charge, so there is an unequal sharing of the + and – charges), then the compound is
polar
 Polar dissolved polar and vice-versa
VSEPR Chart
VSEPR
AB or AB2
AB2E
AB3
AB4
AB3E
AB2E2
AB5
AB6
SHAPE
Linear
Bent
Trigonal-Planar
Tetrahedral
Trigonal-Pyramidal
Bent
Trigonal-Bipyramidal
Octahedral
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The Mole is a counting unit; 6.022 x 1023 molecules, particles or atoms
The atomic mass number, atomic weight/mass or molecular mass is the mass of 1 mole of atoms
(by the way, you do know where to find the molecular mass? Try the periodic chart!)
In Stoichiometry, the key to solving the quandry is the mole
 A balanced chemical equation indicates in simple, small, whole numbers the ratio of compounds
to each other
 If you are given the mass or the volume of only one of the compounds in a chemical reaction,
the remaining information can be calculated or derived using the principles of stoichiometry
 Here’s the deal: in a perfect world and on a perfect planet, compounds will come together an
react in simple whole numbers. We, however, do not exist in said perfect world. Therefore,
we must rely on the simple, small, whole number ratio of compounds
 Example:
6CO2 + 6H2O + energy  C6H12O6 + 6O2
(by the way, this is photosynthesis; look familiar in reverse?)
 Okay, let’s read this reaction. Six moles of carbon dioxide and 6 moles of water and energy
(in the form of light) react to form 1 mole of glucose and 6 moles of oxygen
 The ratio of these compounds is 6:6:1:6
 So, if I start out with 51.4 grams of CO2 divided by the molecular mass which is 44 g/mol I
get a total of 1.17 moles of CO2 to start with.
 So then, the ratio looks something like this: 1.17 : 1.17 : 0.195 : 1.17
 How, you might ask? Well, it’s called a ratio, and we divide. 6:1 is the same thing as
saying 1 : 1/6
 Anyway, that’s how you figure out exactly how many moles of each compound you have.
Then from there, you can multiply by the molecular mass to get actual mass
 Yield is the amount of product that you would expect to recover from a specified amount of
reactant. Theoretical because we perform the experiment on paper first. The limiting factor is
the reactant that theoretically produces the least amount of product.
Reaction types:
 Synthesis: make a new compound from elements or simple substances
 Decomposition: break a complex compound into smaller or elemental parts
 Combustion: a hydrocarbon with oxygen as a fuel burn or combust to form CO2 and H2O
 Single replacement: only one element or polyatomic ion is replaced, exchanged or displaced.
The order of this reaction depends upon the activity chart
 Double replacement: two elements or polyatomic species are replaced; the product must
include a solid; this can be determined by the solubility guidelines
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Molar Heat Capacity (C) is the heat required to raise 1 mol of a substance by 1 degree K or C.
C = q / n . T ; and
o Specific heat (cp) is the heat required to raise 1 gram of a substance by 1 degree K or C.
cp = q / m . T; so,
o C = cp . M (where M is the molar mass in g / mol)
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Thermodynamics is the study of energy and energy is the capacity to do work.
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The 1st Law of Thermodynamics is also the Law of Conservation of Energy: Energy cannot be
created or destroyed, it is conserved in a chemical reaction.
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Enthalpy (H) is the amount of heat absorbed or released in a chemical reaction and is
measured in kJ / mol.
o Enthalpy of formation Hof is the enthalpy to form a compound from its elements at
standard thermodynamic conditions; 25oC and 1 atm.
o Therefore, an element has an enthalpy of formation of “0”.
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Calorimetry is the experimental measure of an enthalpy change and is measured in kJ (calorie
or kilocalorie is another unit)
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In an Exothermic Reaction, the change in enthalpy from the reactants to the product, results in
a negative change in enthalpy. The potential energy for the products is less than the potential
energy for the reactants.
o In this reaction, the energy is on the product side (A + B  C + 22 kJ)
Potential Energy (kJ/mol)
Exothermic Reaction
90
80
70
60
50
40
30
20
10
0
H = Hproducts - Hreactants
Activated Complex
H = C . T
Activation Energy
H =-22 kJ
Energy of Reactants
1
2
3
4
5
6
Energy of Products
7
8
9
10
Reaction Coordinate (time)
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Hess’s Law states that the amount of heat released or absorbed does not depend on the
number of steps; the sum of all changes in enthalpy equals to the net change in enthalpy.
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In an Endothermic Reaction, the change in enthalpy from the reactants to the product, results
in a positive change in enthalpy. The potential energy for the products is greater than the
potential energy for the reactants.
o In this reaction, the energy is on the reactant side (A + B + 30 kJ  C)
Potential Energy (kJ)
Endothermic Reaction
Activation Energy
90
80
70
60
50
40
30
20
10
0
H=30 kJ
Energy of Products
Energy of Reactants
1
2
3
4
5
6
7
8
9
10
Reaction Coordinate (time)
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Entropy (S) is the measure of disorder in a system and is measured in J / K.
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The following events lead to an increase in disorder (entropy):
o Increase in temperature;
o Phase change from solid to liquid to gas;
o More products than reactants;
o Simpler products than reactants;
o Substances that are put into solution;
o Solutions that become more dilute (molecules have more space);
o Gases that the pressure decreases (increases the volume and space).
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S = Sproducts - Sproducts
Molar entropy (So) is the entropy of one mole of a compound.
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Gibb’s Energy (G), also called Gibb’s Free Energy, is the energy in a system available for work
and is measured in kJ / mol.
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The amount of Gibb’s energy predicts a spontaneous reaction which is a reaction that occurs
without any additional energy. A negative G is spontaneous and a positive G is not
spontaneous;
o If,
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H =
+
+
and S =
+
+
-
then
G= and the reaction is
spontaneous
?
spontaneous, low T
?
spontaneous, high T
+
not spontaneous
The factors that affect Gibb’s energy are enthalpy, entropy and temperature
o G = H - TS; and
o G = Gproducts - Greactants
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The 2nd Law of Thermodynamics states that in a spontaneous reaction, entropy is not
conserved and must always increase. All spontaneous reactions contribute a portion of energy
to perform useful work (Gibb’s Energy).
o As a reaction proceeds, G decreases until it equals “0” which is the point of equilibrium.
The states of matter are solid, liquid, gas and sometimes plasma (which is not a naturally
occurring state)
Cohesion is an attraction for particles that a liquid has.
Adhesion is an attractive force for particles of solid surfaces
Capillary action is the motion of a liquid up a small surface and is accomplished by adhesion of
liquid molecules to the surface of the glass as well as cohesion between the liquid molecules.
Surface tension is the fore that acts of the surface of a liquid and tends to minimize the area of the
surface. Why?
 1st of all, cohesive forces bring the molecules of a liquid together so that they stay in contact;
 2nd, under the surface of the liquid, these cohesive forces are pulling equally in all directions;
 3rd, only on the surface, the molecules are being pulled sideways and downward creating
surface tension.
It takes energy to increase the surface area of a liquid because this energy must oppose the net
forces pulling the molecules; conversely, a liquid decrease energy as the surface area decreases.
This tendency toward decreasing the surface area is called surface tension. A high surface tension
means that a lot of energy is needed to break the surface.
Application: surface tension is used in laundry. When you put dirty clothes in the washing
machine and no laundry detergent, the dirt on your clothes cannot penetrate the surface tension
of the water and so it stays on your clothes. When you add the detergent, the soap decreases the
surface tension by disrupting hydrogen bonds and therefore the dirt can be carried away by the
water!!! Although, it also takes effort to load the washing machine!
Endothermic
Reactions
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Gases do not have the same type of intermolecular forces, because they are farther apart and the
attractive forces are minimized. That is why a gas will fill the space available.
Now let’s talk about the physical changes in the states of matter. Refer to the diagram below.
The direction of the arrow indicates the change from one state to another by the addition of
release of energy. As energy is absorbed, the state changes from solid to liquid; liquid to gas; or
solid to gas. As energy is released the state changes in the opposite direction. The name of each
process is labeled.
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Evaporation
Sublimation
Melting
Gas
Condensation
Liquid Deposition
Freezing
Solid
Exothermic
Reactions
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The process by which a solid changes into a liquid by the absorption of energy or the decrease in
pressure is melting and the temperature that this occurs is the melting point. At this temperature,
matter has absorbed enough energy for the molecules to break free from the attractive forces of
the solid.
The process by which a liquid changes into a solid by the release of energy or the increase in
pressure is freezing and the temperature is called the freezing point. At this temperature, the
particles of a molecule are slowed to a point where the intermolecular forces of attraction hold
them in a crystalline state.
Covalent bond types are a result of intermolecular forces (or van der Waals forces) or the forces of
attraction that hold molecules together. This is a review. See the chart below for a summary:
Bond Type
Polar Covalent
Bonds
Hydrogen Bonds
Non-Polar
Covalent Bonds
Ionic Bonds
Intermolecular Forces
(covalent)
Dipole – Dipole Forces
Explanation of Forces
Positive and negative ends attract each other.
The greater the electronegativity difference,
the more polar the bond is. The more polar a
bond, the stronger the attractive forces,
stronger the bond.
Stronger dipole forces
Hydrogen bonds with very electronegative
between neighboring
atoms as Oxygen, Nitrogen or Fluorine of
molecules
another molecule and leaves the Hydrogen
with a large partial positive charge. Because
hydrogen only has one electron, which leaves
behind a very strong proton
London Dispersion Forces Weak attraction caused by temporary dipoles
from the uneven distribution of electrons;
interactions occur between the negative part
of one molecule and the positive region of a
neighboring molecule.
Electrostatic Forces
Cation and anion attractions
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•
•
•
•
•
•
•
•
•
•
•
•
The stronger the intermolecular forces and the stronger the bond, the greater the energy needed
to break the bond. Molecules with stronger bonds have higher boiling and melting points and tend
to be solids or liquids.
Water has the ability to form two hydrogen bonds per one molecule of water which leads to
greater versatility.
London dispersion forces increase with increasing molecular mass and decrease with increasing
distance between nuclei; in other words, the melting and boiling points increase with increase in
mass.
ORDER OF STRENGTH OF INTERMOLECULAR FORCES
hydrogen bonding – strongest
dipole – ion or electrostatic forces in ionic bonding
dipole – dipole
London dispersion forces
Nonpolar substances are usually gases at room temperature or have low boiling points because of
the low intermolecular forces
Polar substances have high boiling points – many are solids at room temp. (ionic compound –
strong intermolecular forces)
Properties affected by intermolecular forces
Boiling point
Retention of volume and shape
Surface tension
Evaporation
Vapor pressure
Viscosity
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The enthalpy of fusion (symbol: ΔfusH), also known as the heat of fusion, is the amount of
thermal energy which must be absorbed or evolved for 1 mole of a substance to change states
from a solid to a liquid or vice versa. It is also called the latent heat of fusion or the enthalpy
change of fusion, and the temperature at which it occurs is called the melting point.
 Hfus = H(liquid at melting point) – H(solid at melting point)
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When you withdraw thermal energy from a liquid or solid, the temperature falls. When you add
heat energy the temperature rises. However, at the transition point between solid and liquid (the
melting point), extra energy is required (the heat of fusion). To go from liquid to solid, the
molecules of a substance must become more ordered. For them to maintain the order of a solid,
extra heat must be withdrawn. In the other direction, to create the disorder from the solid crystal
to liquid, extra heat must be added.
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The standard enthalpy change of vaporization, ΔvHo, also (less correctly) known as the heat
of vaporization is the energy required to transform a given quantity of a substance into a gas. It is
measured at the boiling point of the substance, although tabulated values are usually corrected to
298 K: the correction is small, and is often smaller than the uncertainty in the measured value.
Values are usually quoted in kJ/mol, although kJ/kg, kcal/mol, cal/g and Btu/lb (obsolete) are also
possible, among others.
 Hvap = H(vapor at boiling point) – H(liquid at boiling point)
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Molar entropy of vaporization, Svap are for the gaseous states.
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The standard enthalpy change of condensation (or heat of condensation) is numerically
exactly equal to the standard enthalpy change of vaporization, but has the opposite sign.
Enthalpy changes of vaporization are always positive (heat is absorbed by the substance), whereas
enthalpy changes of condensation are always negative (heat is released by the substance).
Melting point = Tmp = Hfus / Sfus and;
Boiling point = Tbp = Hvap / Svap
A phase is a uniform collection of particles.
Some things can exist in two phases (milkshakes, slushies, etc.)
Equilibrium is where G = 0; forward and reverse reactions occur at the same rate and the
concentration of each stays the same.
Dynamic equilibrium is a state of a compound where the particles move between 2 different states.
Vapor pressure is the pressure produced by a liquid or a solid when it is in dynamic equilibrium
with its gas phase and is measured in mmHg or kPa.
As temperature increases, vapor pressure increases and exerts pressure on the walls of the
container.
PA phase diagram shows the state of a compound with temperature and pressure.
The triple point is the temperature and pressure where the three phases exist in equilibrium
The critical point (critical temperature and critical pressure) is the point above which the liquid and
gas phases are indistinguishable (a supercritical fluid). Above the critical temperature, a gas
cannot be liquefied.
As vapor pressure increases, density increases.
Pressure plays a greater role in the phases of liquids and gases than solids.
To draw a phase diagram, plot the critical point, the triple point and the mp/bp.
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Solution: a homogenous mixture of two or more substances in a single phase; a solute and a
solvent.
Solute: a substance which is the portion of a solution that is dissolved.
Solvent: a substance which is the portion of a solution that dissolves the solute; determines
the phase of the solution.
How does a solvent dissolve the solute, you might ask? The solvent has enough energy to break
the surface tension or the bonds of the solute. Remember those pesky intermolecular forces?! The
strongest of all are the metallic bonds followed by the ion-dipole, hydrogen, dipole-dipole, and
finally the London dispersion forces.
Table 2: Phases of solutions.
Phase of the
Phase of the
Solute
Solvent
Solid
Solid
Solid
Liquid
Phase of the
Solution
Solid
Liquid
Example
Metal alloy
Kool-aid
Liquid
Gas
Gas
Liquid
Liquid
Gas
Liquid
Liquid
Gas
Ethanol-water
Dr. Pepper
Air
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Suspension: A mixture whose particles are evenly dispersed in a gas or a liquid and which settle
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Colloid: a mixture whose particles are smaller than a suspension but larger than a solution and
does not settle; the particles are suspended in a liquid, gas or solid. The Tyndall effect is a
out over time.
Examples: sand – water mixture, cement
phenomenon where a beam of light is reflected off of the particles and visible to the naked eye.
Examples: milk, mayonnaise, marshmallow, fog
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One difference in mixtures and compounds is that mixtures can be separated into it’s parts.
Mixtures can be separated by chromatography, gel electrophoresis, distillation, mechanical
separation, evaporation, etc. Can you think of any more?
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Concentration is the measure of the amount of a particular substance in a given volume of
solution.
Table 3: Concentration calculations
Name
Units
Application
Molarity
M
mol solute / L solution
Solution concentration
Molality
m
mol solute / kg solvent
Calculations using solids
Normality
N
Molarity x n
Titration calculations
Parts per Million
ppm
g solute / 1 000 000 g solution
To express very small
6
mass solute / mass solution x 10
concentrations
Percentage by Weight
%
Mass solute/ mass solution x 100
% by weight
Percentage by Vol.
%
Vol. solute / vol. solution x 100
% by volume
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Molarity: The molar unit is probably the most commonly used chemical unit of measurement.
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Molality: The molal unit is not used nearly as frequently as the molar unit. A molality is the
Molarity is the number of moles of a solute dissolved in a liter of solution. A molar
solution of sodium chloride is made by placing 1 mole of a solute into a 1-liter volumetric flask.
Water is then added to the volumetric flask up to the one liter line. The result is a one molar
solution of sodium chloride.
M = n/V
(remember that n = mass / molar mass)
M = 1 mol NaCl / 1 L H2O
&
1 mol NaCl = x grams NaCl / 58.44 g/mol NaCl
So,
M = 58.44 grams NaCl / 1 L H2O
number of moles of solute dissolved in one kilogram of solvent. Be careful not to confuse
molality and molarity. Molality is represented by a small "m," whereas molarity is represented by
an upper case "M." Note that the solvent must be weighed unless it is water. One liter of water
has a specific gravity of 1.0 and weighs one kilogram; so one can measure out one liter of water
and add the solute to it. Most other solvents have a specific gravity greater than or less than one.
Therefore, one liter of anything other than water is not likely to occupy a liter of space. To make a
one molal aqueous (water) solution of sodium chloride (NaCl) , measure out one kilogram of water
and add one mole of the solute, NaCl to it. The formula weight for NaCl is 58, and so 58.44 grams
of NaCl dissolved in 1kg water would result in a 1 molal solution of NaCl.
m = 1 mol NaCl / 1 kg H2O = 58.44 g NaCl / 1 kg H2O
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Percent by weight: To make up a solution based on percentage by weight, one would simply
determine what percentage was desired (for example, a 20% by weight aqueous solution of
sodium chloride) and the total quantity to be prepared.
If the total quantity needed is 1 kg, then it would simply be a matter of calculating 20% of 1 kg which, of
course is: 0.20 NaCl * 1000 g/kg = 200 g NaCl/kg.
In order to bring the total quantity to 1 kg, it would be necessary to add 800g water.
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Making dilutions is easy. The starting solution is referred to as the stock solution and the solution
that you make is referred to as the working solution. It is easier to make and store smaller
volumes of concentrated stock solution. When you need a reagent (A substance used in a
chemical reaction to detect, measure, examine, or produce other substances;) simply prepare it
from a stock solution. The molarity and volume of the stock solution are inversely proportional.
As you reduce one, you increase the other by the same proportion. Therefore;
MstockVstock = MworkingVworking
For example: You start out with a 12M concentrated solution of saline solution. The experiment
calls for 0.8 Liters of a 3.0M solution. How would you prepare it?
12M x Vol. Stock = 3.0M x 0.8L
Vol. Stock = 3.0M x 0.8L / 12M = 0.2L
Therefore, you take 0.2L of your stock solution and dilute to 0.8L final volume.
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What makes a substance soluble? Solubility is the relative ability of a substance to dissolve into
another at a given temperature and pressure; or the amount of solute that will dissolve into a
given solution to saturate the solution. A saturated solution cannot dissolve any more solute –
like a glass of iced tea that you added too much sugar to and the sugar settles on the bottom. An
unsaturated solution contains less solute that the saturated solution – like the perfect glass of
sweetened iced tea, with no settled solute (sugar) on the bottom. A supersaturated solution
holds more dissolved solute than what is required to reach solubility equilibrium at a given
temperature - or a solution that contains more solute than it would if the dissolved solute were in
equilibrium with the undissolved solute. A solubility equilibrium is the state at which a solution
is in equilibrium with its solute ions and the solvent.
 SO, try this at home. Saturate a pan of water with sugar, add more sugar so that it settles on
the bottle. Heat the water. What happens to the excess sugar? (It dissolves into the water.)
Allow the water to cool. Did the excess sugar fall out of solution? (No) Now, that is a
supersaturated solution!
The speed of solubility depends on the intermolecular forces, mass, surface area, temperature. As
temperature increases, solubility increases.
What effects do a solute have on a solvent? To answer that question, consider what happens
when you apply salt (NaCl) to an icy road. Does the salt melt the ice? No, what the salt does is
lower the melting or freezing point of water . What happens when you add salt to a pot of water
on the stove? It raises the boiling point of the water. So using this example, a solute (the salt)
decreases the melting/freezing point of the solvent (water) and raises the boiling point. The
melting/freezing and boiling points of a solution are colligative properties. A colligative property
is a property of a solution that is dependent upon the number of solute particles present and the
nature of the solvent. These properties are independent of the identity of the solute. In other
words, if you replaced 1 mole salt with 1 mole of sugar in the previous example, the effects on
mp/fp and bp would be the same! The greater the concentration of solute, the greater the effect
it has on a colligative property. A solute also lowers the vapor pressure of a given solvent.
 The change in the bp and mp can be calculated for solutions if the solute is molecular (covalent
bonding, not ionic) and non-volatile (does not evaporate) because the bp and fp are
proportional to the molality.
 Tb = (+kb) x (m) and; Tf = (-kf) x (m) where “+” indicates and increase and “-“ indicates a
decrease in temperature. The symbols kb and kf are the boiling point and freezing point
constants (oC/mol) and represents the number of degrees in Centigrade that the bp or fp is
raised or lowered when 1 mole of a molecular, non-volatile solute is dissolved in 1 kg of a
solvent.
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Henry’s Law states that at a constant temperature, the solubility of a gas is directly proportional
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Sometimes, solutions exhibit different characteristics than the parent solute or solvent. (Does that
sound familiar? How about mother and daughter ions!) An electrolyte is a solute that is able to
conduct electricity when dissolved in a solvent. It is said to have conductivity. Conversely, a
non-electrolyte does not conduct electricity in solution or out of solution. Ions can be produced
in solution 2 ways:
1. An ionic compound is separated into it’s ions by a process called dissociation; or
2. A polar covalent compound dissolves in water and loses a Hydrogen to form a Hydronium ion:
HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq)
The effect of these electrolytes on the bp and fp are higher or lower than expected based on the k
x m equation. Why? Because, 1 mole of an ionic compound produces more than 1 mole of
electrolytes! Let’s see this in action. One mole of NaCl produces 1 mole of Na+ and 1 mole of Clfor a total of 2 moles of electrolytes. Therefore, the effect on the bp/fp is 2 times the effects of a
molecular substance.
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to the partial pressure of the gas on the surface of the liquid. When you open a soda, either in a
bottle or in a can, there is a sudden release of pressure and the gas dissolved on the surface of
the soda, comes out of solution (it foams up.) Soda is a saturated solution of CO 2 and water,
sugar, etc. The carbon dioxide is exerting partial pressure on the surface of the soda. Soda has a
high pressure when it is first bottled or canned. This high pressure environment increases the
solubility of the CO2; therefore, more carbon dioxide is dissolved into the soda. When the pressure
is released, when the soda is opened, the pressure in the container equals the pressure of the
atmosphere, forcing the excess CO2 to come out of solution. (I can’t wait to talk about equilibrium
constants!)
What do you wash your clothes with? In Chemistry terms, laundry detergent is referred to as a
surfactant. A surfactant or surface active agent, is a molecule that is amphiphilic which means
that it has a water loving end (hydrophilic) and water fearing end (hydrophobic). These agents are
wetting agents that lower the surface tension of a liquid, allowing easier spreading, and lower the
interfacial tension between two liquids. Detergent is a manufactured water soluble surfactant that
emulsifies oil and dirt; makes a suspension from one liquid into another where the first liquid does
not dissolve in the second (kind of like oil and vinegar salad dressing). Soap is a naturally
occurring detergent that emulsifies dirt.
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Acids and bases are electrolytes
In solution acids are sour and bases are bitter
Vinegar is an acid; Alum (which is used to make pickles) is a base
Acids react with certain metals to make hydrogen gas
Acids and bases neutralize each other
In other words, they raise/lower the pH to a non-acid/non-base
Acids and Bases can both destroy human tissue
Naming Binary Acids
o Acids are formed when a Hydrogen is added to an anion
o Take the anion root,
o Add a hydro- in front
o Change the ending from –ide or -ate to –ic; or from –ite to -ous
o Put them all together and add the word “acid”
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Naming Bases; Simply name the cation and anion
There are 3 different types of acids and bases
o Arrhenius
o Brönsted-Lowry
o Lewis
Arrhenius Acids and Bases
o An Arrhenius Acid is a substance that releases H+ ions in an aqueous solution
o An Arrhenius Base is a substance that releases –OH in an aqueous solution
Brönsted-Lowry Acids and Bases
o A Brönsted-Lowry acid is any species that can donate a proton (H+ ion) to another
species; a proton donor
o A Brönsted-Lowry base is any speices that can accept a proton (H+ ion) from another
species; a proton acceptor
Amphiprotic: A compound that can be both a proton donor OR a proton acceptor in separate
reactions is called amphiprotic or amphoteric
Lewis Acids and Bases
o A Lewis Acid is any species that can accept a pair of electrons from another species;
electron pair acceptor
o A Lewis Base is any species that can donate a pair of electrons from another species;
electron pair donor
Acid-Base Equilibria
In a chemical reaction involving both acids and bases, dynamic equilibrium can still be achieved
The strength of an acid depends on a number of factors, such as the properties of the solvent,
the temperature, and the molecular structure of the acid.
We compare the strengths of two acids, in the same solvent and at the same temperature.
That way we can focus on the structure of the acid.
Ionization Constants of Acids and Bases
o HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq)
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o Keq= [H3O+][A-] / [HA][H2O]
o Keq= [H3O+][A-] / [HA]
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pH + pOH = 14
o pH = -log10[H3O+]
o pOH = -log10[OH-]
Titration is a common laboratory method of quantitative/chemical analysis which can be used
to determine the concentration of a known reactant
o MV = MV
o The titrant is usually the solution of known concentration that is delivered by a burette
into a known quantity of the solution of unknown concentration
o The titraver is the indicating solution
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Ionization Constants
o Kw= [H3O+][OH-]
o Kw = 1.0 x 10-14
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Oxidation-Reduction: “Redox Reactions”
Oxidation  Loss of electrons
 Na  Na+ + eReduction  Gain of electrons
 Cl + e-  ClOxidation Number (Oxidation State) 
 The charge that an atom would have if the electrons in the bond were possessed
entirely by the more electronegative element.
Oxidation numbers serve as a bookkeeping tool used to keep track of electron movement.
Assigning Oxidation Numbers
o All pure elements and homogeneous molecules = 0
o Elements in group IA = +1
o Elements in group IIA = +2
o Ag+, Zn+2, Al+3
o In binary compounds the second element = anion charge
o Oxygen is almost always = -2
o Hydrogen is almost always = +1
o The total charge of a compound is always = 0
Oxidizing/Reducing Agents
To determine if a reaction is redox, determine the oxidation numbers for all elements.
o The element that was oxidized is part of the reducing agent.
o The element that was reduced is part of the oxidizing agent.
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