300 Chemistry Finals Review – TOPICS:
• Limiting Reactants (continuation of stoich)
• Gases
– Gas Laws and Kinetic Molecular Theory
• Nuclear
– Radioactive decay; half-life; energy
transformation
• Bonding
– Lewis Structures, VSEPR and polarity
• States of Matter
– IMFs, KE and PE, heating/cooling curves
300 Chemistry Finals Review
• Solutions
– Molarity, solubility, colligative properties
• Thermochem., Kinetics & Equilibrium
– PE diagrams, Collision Theory, Keq
• Acids and Bases
– Properties, Kw, pH loop, neutralization,
titration, Arrhenius, Bronsted-Lowry
Limiting Reactants
• In a chemical reaction, one reactant is
always used up first:
– Limiting Reactant (used up completely)
– Excess Reactant (some left over)
• Stoichiometry calculation adjustments:
– Once limiting reactant is used up, reaction
STOPS! No more product is made.
– Determine how much product each reactant
could theoretically make. The one that
actually yields the LEAST product is “limiting”
– % yield = Actual / Theoretical Yield x 100
•States of Matter
• What makes a gas? How is this “state of
matter” different than solids or liquids?
• Why are some substances gases and others
solids at the same temperature?
• Differing attractive forces between
molecules cause some materials to be
solids, some to be liquids, and some to be
gases at the same temperature.
• In gases, particles have enough energy of
motion that that they are no longer stuck
together and are free to move around until
they hit the walls of their container.
What is Gas Pressure?
1 atm =
760 mmHg
•
PRESSURE is a force exerted by
the substance per unit area on
another substance.
•
GAS PRESSURE is the force that
the gas exerts on the walls of its
container.
•
A balloon expands because the
pressure of the gas molecules is
greater than the pressure of the
gas molecules on the outside.
•http://www.indiana.edu/~geog109/topics/10_Forces&Winds/GasPressWeb/PressGasLaws.html
The Combined Gas Law (cont.)
The Combined Gas Law
• The combined gas law states the
relationship among pressure, temperature,
and volume of a fixed amount of gas.
Avogadro's Principle (cont.)
• Recall that one mole of a substance
contains 6.022 x 1023 particles (number is
named after Avogadro).
• The molar volume of a gas is the volume
1 mole occupies at 0.00°C and 1.00 atm of
pressure.
• 0.00°C and 1.00 atm are called Standard
Temperature and Pressure (“STP”).
• At STP, 1 mole of ANY gas occupies 22.4 L.
The Ideal Gas Law (cont.)
• The ideal gas constant is represented by R
and is measured to be 0.0821 L•atm/mol•K
when pressure is in atmospheres.
• The ideal gas law describes the physical
behavior of an ideal gas in terms of pressure,
volume, temperature, and amount (n=moles)
PV = nRT
The Ideal Gas
• The ideal gas law: PV = nRT
If pressure is 2.00 atm, volume is 17.0 L
and the temperature is 298 K, how many
moles of gas do you have?
n = PV / RT
= (2 atm x 17.0L) / (0.0821 L*atm/mol*k x 298K)
= 1.39 moles
Stoichiometry of Reactions Involving Gases
• The gas laws can be applied to calculate
the stoichiometry of reactions in which
gases are reactants or products.
2H2(g) + O2(g) → 2H2O(g)
• 2 mol H2 reacts with 1 mol O2 to produce
2 mol water vapor.
• At a constant P and T, the ratios of moles and
volume are the same!
• Thus, 2 L H2 would fully react with 1 L O2 to
produce 2 L water vapor.
Section Assessment
How many mol of hydrogen gas are
required to react with 1.50 mol oxygen gas
in the following reaction?
2H2(g) + O2(g) → 2H2O(g)
A. 1.00
B. 2.00
C. 3.00
D. 4.00
A.
B.
C.
D.
A
B
C
D
Section Assessment
How many liters of hydrogen gas are
required to react with 3.25 liters of oxygen
gas in the following reaction?
2H2(g) + O2(g) → 2H2O(g)
A. 2.00
B. 3.25
C. 4.00
D. 6.50
A.
B.
C.
D.
A
B
C
D
Gas Pressure (cont.)
• Dalton’s law of partial pressures states
that the total pressure of a mixture of gases
is equal to the sum of the pressures of all
the gases of the mixture.
Ptotal = P1 + P2 + P3 +...Pn
• Remove water vapor pressure when a
gas has been “collected over water’. Can
look up the Pwater in a table and subtract
from total pressure based on Dalton’s Law.
Nuclear Change
Nuclear change:
Definition: when an unstable nucleus reacts to
form a more stable nucleus
Significant ENERGY changes occur during nuclear
reactions compared to chemical reactions.
Nuclear reactions involve the release of charged
particles, electromagnetic waves (“rays”) or both.
This process is called:
RADIOACTIVITY!
Types of Nuclear Reactions
• Decay = Unstable radioisotopes of an element
spontaneously break down into stable (nonradioactive)
isotopes of another element by emitting particles
• Capture = Particles hit a nucleus and are attached, causing
the atom to change into another element
• Bombardment = Particles strike a nucleus and split it,
causing it to change into another element
• Fission = A large nucleus is bombarded with neutrons and
splits into smaller fragments, releasing energy
• Fusion = Small nuclei combine at very high temperatures to
form a larger nucleus and release energy
Nuclear Change
Types of radiation:
– Beta Decay (β particle) – negatively charged
particle (electron) emitted during a certain type of
radioactive decay, known as beta decay
– Beta Decay stabilizes nuclei by converting
neutrons to protons
n
Important
Point!
p + e-
In Beta decay, a neutron emits a
high-energy electron (a beta
particle) and changes into a
proton!
Section 19.1
Radioactivity
A. Radioactive Decay
Types of Radioactive Decay
• Beta-particle production
• Beta particle – electron
– Examples
• Net effect is to change a neutron to a proton.
Nuclear Change
Alpha Decay
Unstable nucleus can get more stable by
emitting a positive “alpha particle”
Alpha particle (α) = 2p+ + 2no = helium nucleus
 2 protons and 2 neutrons (mass number = 4)
From: http://library.thinkquest.org/3471/radiation_types_body.html
Section 19.1
Radioactivity
A. Radioactive Decay
Types of Radioactive Decay
• Alpha-particle production
• Alpha particle – helium nucleus
– Examples
• Net effect is loss of 4 in mass number and loss of 2 in
atomic number.
How Penetrating are these Rays?
• Alpha particles
– Large, charged and slow moving. Easiest to stop. (Skin or paper)
• Beta particles and Positrons
– Smaller and faster, but still charged, so they stick to other particles.
– Can be stopped by thin Al foil
• Neutrons
– Large, uncharged particles – highly penetrating.
– Need thick concrete to stop.
• Gamma and X-rays
– High energy (fast), massless, electromagnetic waves.
– Very highly penetrating. Stop only with thick lead shield.
Fission and Fusion
• Fission reaction: an atom absorbs a
neutron which splits it into pieces and a
large amount of energy is given off as
heat
Fission and Fusion
• Fusion: two light nuclei (usually H) join
together to form a heavier nucleus and
give off energy.
Energy Transformations:
Nuclear Power Generator
(http://www.nrc.gov/reading-rm/basic-ref/students/animated-pwr.html)
Section 19.1
Radioactivity
C. Detection of Radioactivity and the Concept of Halflife
• Half-life – time required for
half of the original sample
of radioactive nuclides to
decay
• Half-life problems require
you to solve for:
• Half-life time
• Fraction remaining
• Time elapsed
• Initial amount
Section 19.1
Radioactivity
Decay of a Radioactive Element
Half of the
radioactive
parent atoms
decay after one
half-life. Half of
the remainder
decay after
another half-life
and so on……..
Half-life activity
Covalent Bonds
• Bonding occurs because the atoms are more stable
state afterwards. The total energy of the atoms is
lower after bonding.
• Covalent bonds are formed when atoms share
valence electrons. Caused by the attraction of
protons in one atom for electrons in another.
• Sharing valence electrons with other atoms results
in pseudo noble-gas electron configurations (with
filled valence shells; octet rule obeyed).
• A covalent molecule is formed when two or more
non-metal atoms bond into a single unit.
Fluorine’s Covalent Bond
(F2 molecule)
Creating Octets for Everyone!
Covalent Bonds
form when
unpaired electron
orbitals overlap.
Polar Bonds (dipoles)
• Some covalent bonds have a “+” and “–” side
and are called “dipoles” or “polar bonds”.
• Electronegativity differences between atoms
causes the shared electron to spend more
time near the more electronegative element.
Example: HBr
+
-
Dipoles (Polar Bonds)
• Electronegativity differences creates “polar”
bonds with partial +/- charges on each side.
+
-
• “Dipole moment” is drawn as an arrow from
the positive to the negative centers of charge.
Bond Character (% ionic)
Electronegativity Difference = Type of Bond
• Nonpolar covalent bonds form between atoms
of similar electronegativities (<0.4).
– Between atoms of same nonmetal element.
• Polar covalent bonds form between atoms of
moderately different electronegativities (0.4-1.7)
– Between atoms of different nonmetal elements.
• Ionic bonds form between atoms of highly
different electronegativities (>1.7).
– Between atoms of metals and non-metal elements.
Multiple Covalent Bonds
• Double bonds form when two pairs of
electrons are shared between two atoms.
• Triple bonds form when three pairs of electrons
are shared between two atoms.
.
“VSEPR” Model
• The Valence Shell Electron Pair Repulsion
model helps predict the 3D geometries of
different molecules.
• It is based on the premise that electrons
around a central atom repel each other, so
the atoms spread out as far apart as possible.
• Lone pairs of electrons are more repulsive
than bonding pairs.
Lone-pairs
Central Atom
Bonding-pairs
“VSEPR” Molecule Geometries
• 3D Molecular shapes and their “bond angles”:
Linear
Trigonal planar
Bonding-pairs
Lone-pairs
Tetrahedral
Pyramidal
Bent
Molecular Polarity
• Recall that electronegativity differences creates
“polar” bonds between different non-metal atoms
• “Dipole moment” is drawn as an arrow from the
positive to the negative centers of charge.
• How do we determine if the overall molecule is polar?
(i.e. that it has a positive and negative side)
Molecular Polarity
• If all the bonds in a molecule are NON-polar:
– The entire molecule is non-polar!
– This always occurs in molecules made of one element.
• A molecule will be POLAR only if:
– There are polar bonds present
– The molecule is “asymmetrical” (lopsided) either in
geometry or in its terminal elements (i.e. different elements)
• Symmetrical 3D geometries have NO lone pairs:
– linear, trigonal planar, tetrahedral
• Asymmetrical 3D geometries have lone pairs:
– Bent, pyramidal
Molecular Polarity
Examples:
Hydrogen Bromide (HBr) =
Polar Molecule (diff elements)
BF3 = Non-Polar Molecule
(Symmetrical geometry)
Ammonia (NH3) =
Polar Molecule
(Asymmetrical geometry)
Molecular Polarity Method
1. Draw the Lewis Structure for molecule
2. Determine if bonds are polar or non-polar.
• If no bonds are polar, then molecule is non-polar (done!)
3. If polar bonds exist, determine VSEPR geometry:
4. If geometry is asymmetrical, then polar molecule
5. If symmetrical – check to see if terminal atoms are all
the same element (i.e. same electronegativity)
• If elements are different, then molecule is polar
• If elements are all the same, then molecule is non-polar
(because they are all pulling electrons equally in all directions)
Polarity “Decision Tree”
Non - Polar
Terminal
Atoms?
Polar
Geometry?
Type of
Bonds?
Polar
Non - Polar
Intramolecular Forces
•
Relative strength of attractive forces cause some materials to be solids,
some to be liquids, and some to be gases at the same temperature.
Intermolecular Forces
•
IMFs are much weaker than covalent bonds (10%
or less – often less than 1% as strong).
•
Three main IMFs:
• Dipole-Dipole
• Hydrogen Bonding
• London Dispersion Forces
Intermolecular Forces
•
•
Dipole-dipole forces are attractions between
oppositely charged regions of polar molecules.
Much weaker than covalent bonds.
Intermolecular Forces
•
•
Hydrogen bonds are very strong, special dipole-dipole
attractions that occur between hydrogen atoms in a polar
bond and a small, highly electronegative atom, typically
F, O or N (“phone”).
Water molecules display hydrogen bonding, which
explains its unique properties (high boiling point, etc).
Intermolecular Forces
• London Dispersion forces are weak forces that result
from temporary shifts in density of electrons in
electron clouds. These are the only IMFs in nonpolar
molecules, but also exist in polar molecules as well.
• The more electrons there are in the molecule, the
greater the London dispersion force.
Intermolecular Forces
•
What is the dominant IMF for each of these molecules:
•
O2
= London Dispersion forces (non-polar mlc)
•
HBr
= dipole-dipole (polar mlc)
•
H2O
= hydrogen bonding (polar mlc with H + O)
•
CH4
= London Dispersion forces (non-polar mlc)
•
C2H5OH = hydrogen bonding (polar mlc with H + O)
•
NH3
= hydrogen bonding (polar mlc with H + N)
•
SiOBr2
= dipole-dipole (polar mlc)
•States of Matter (Phases)
State
Solid
Shape
Volume
Definite
Definite
Particle
Arrangement
and
Compressibility
Closely packed,
not compressible
Particle
Movement
Speed of
Particle
Movement
Kinetic
Energy of
Particles
None, slow
vibrations
in place
None/slow
vibrations
Very low
Liquid No
Definite
More loosely
packed, only
slightly
compressible
Particles
can slide
past each
other
Moderate
Low to
moderate
No
No
definite
Very far apart,
very
compressible
Fast,
random
movement
Fast
Very high
Gas
•Phases & Intermolecular Forces
• Key Concept Alert!!!
• Temperature is a measurement of the
average kinetic energy (motion) of the
particles of a substance.
• All substances at the same temperature,
whether they are solid, liquid or gas will have
the same kinetic energy.
• The reason different substances can exist in
different states of matter at the same
temperature is that the have different forces
of attraction holding them together.
Phases & Kinetic Molecular Theory
• The state of matter of a substance is based on the
relative strength of IMFs vs. kinetic energy of
molecules.
• Solids: when Kinetic Energy << intermolecular
forces, the molecules will stay firmly attached and
don’t move relative to each other.
• Liquids: when Kinetic Energy < or = IMFs, the
molecules will stick to each other, but can move
around relative to each other (flow)
• Gases: when KE > IMFs, particles can break free!
No sticky together at all. Free to move anywhere.
Phase Changes
• Phase changes that require energy to occur:
 Melting
 Vaporization
 Sublimation
• Phase changes that release energy:
 Freezing
 Condensation
 Deposition
•Phase Changes Involve Heat
Transfer
• Heat is the transfer of energy from an
object at a higher temperature to an
object at a lower temperature.
•Phase Changes That Require Energy (cont.)
• The boiling point is the temperature at
which the vapor pressure of a liquid equals
the atmospheric pressure.
Heating Curves - Where is the energy going?
Kinetic energy increasing
•
Heat of vaporization:
Potential energy
increasing 
Kinetic energy increasing
•
Heat of fusion
Kinetic energy increasing 
 Potential energy increasing
•Vapor Pressure Curves
• Vapor pressures increase with increasing
temperature since the molecules are gaining
kinetic energy.
• Each substance has its own vapor pressure
curve. Based on the strength of their attractive
forces (IMF’s, etc.).
• The stronger the attractive forces, the lower the
vapor pressure at a temperature.
Homogeneous Mixtures = Solutions
• Solutions are homogeneous mixtures that
contain two or more substances called the
solute and solvent.
• Term is typically used with liquids (water),
but gases and solids form solutions too.
• The solvent is the most abundant material
in the mixture. Water is our key solvent.
• Solutes are the less abundant material that
is mixed (“dissolved”) in the solvent.
• Can be liquids, solids or gases…all dissolve!
Solutions
• Electrolytes
– Compounds that dissociate into separate ions
in water and are good conductors.
– “Strong” electrolytes = 100% dissociation
• Ionic compounds are strong electrolytes
• Also includes the strong acids: HCl, HNO3, HBr
• Fully dissociate to form H+ ions in solution (H3O+)
– “Weak” electrolytes < 100% dissociation
• Nonelectrolytes
– Do not dissociate into ions in water.
– Nonpolar molecules
Expressing Concentration
(cont.)
• Molarity is the number of moles of
solute dissolved per liter of solution.
• Dilution equation: M1V1 = M2V2
Expressing Concentration
(cont.)
• Molality is the ratio of moles of solute
dissolved in 1 kg of solvent.
The Solvation (Dissolving) Process
• Solvation is the process of surrounding solute
particles with solvent particles to form a solution.
• LIKE DISSOLVES LIKE!
• Polar water molecules dissolve polar solutes
(ionic compounds and polar molecules).
• Nonpolar solvents dissolve nonpolar solutes.
• Solvation in water is called hydration.
• The water molecules surround the solute particles.
• Dipole-ion intermolecular force: Attraction between
dipoles of a water molecule and ions of a crystal are
greater than attraction among ions of a crystal.
Solubility
• Solubility = the amount a solute will dissolve
in a solvent. It depends on the type of solute
and solvent and temperature.
• As concentration of solute in solvent increases,
more solute particles collide with remaining
crystalline solid and precipitate out of solution.
• Saturated solutions contain the maximum
amount of dissolved solute for a given amount
of solute at a specific temperature and
pressure.
• Equilibrium is reached between solvation and
precipitation.
Factors That Affect Solvation
• Agitation: Stirring or shaking moves
dissolved particles away from the contact
surfaces more quickly and allows new
collisions to occur.
• Breaking the solute into small pieces
increases surface area and allows more
collisions to occur.
• As temperature increases, rate of solvation
increases.
• Solubility is affected by increasing the
temperature of the solvent because the kinetic
energy of the particles increases.
Temperature
can make a
big
difference!
Colligative Properties of Solutions
• Colligative properties are physical properties
of solutions that are affected by the number of
particles but not by the identity of dissolved
solute particles.
• Types of colligative properties:
• Vapor pressure lowering
• Boiling point elevation
• Freezing point depression
• You depend upon two of these properties every
time you drive your car.
Colligative Properties of Solutions
• The relationship between moles of solute and moles of
particles in solution is the van’t Hoff factor:
“ i ” = (moles of particles) / (moles of solute)
• Note: Molecules always have an i = 1
• Strong electrolytes have i = 2 or more depending upon
the number of ions in formula (CaCl2 has i = 3).
Boiling Point Elevation (cont.)
• The temperature difference between a solution’s
boiling point and a pure solvent's boiling point is
called the boiling point elevation.
ΔTb = iKbm
• where
• ΔTb is the “boiling point elevation”,
• i is the van’t Hoff factor,
• Kb is the molal boiling point elevation
constant, and
• m represents molality.
Freezing Point Depression (cont.)
• Solute particles interfere with the attractive
forces among solvent particles.
• A solution's freezing point depression is the
difference in temperature between its new
freezing point and the freezing point of the
pure solvent.
ΔTf = iKfm
• where ΔTf is the freezing point depression, i is
the van’t Hoff factor, Kf is the freezing point
depression constant, and m is molality.
Equilibrium

Different types of chemical reactions:

One-Way reactions (single arrow pointing right)
We have been assuming that they “go to completion”



Burning wood (glucose):

C6H12O6 + 6 O2  6 CO2 + 6 H2O + energy

Reversing this is hard = energy and time and multiple steps
Reversible reactions (two way arrows)
–
The reaction can occur in BOTH directions
–
Usually simple, single step reactions
 N2(g) + O2(g) ⇄ 2 NO(g)
Equilibrium

For chemical reactions, equilibrium is reached
when the Change in CONCENTRATIONS of
reactants and products stops
Equilibrium


Why do reactions reach Equilibrium?
–
“Effective collisions” concept helps explain:
–
As concentration of “reactants” declines, they get
fewer opportunities for effective collisions, thus the
rate of the forward reaction slows.
–
As concentration of “products” increases, there are
more chances for effective collision in reverse
direction (the “products” become reactants!).
At equilibrium:
Rate of Forward Reaction = Rate of Reverse Reaction
A System at Equilibrium
As rates become equal
…Concentrations become constant
Direction and Sign of Heat Flow
Heat Flow = q
“q” is + when heat flows into the system from the
surroundings
“q” is - when heat flows out of the system into the
surroundings
An endothermic process (q > 0) is one in which heat
flows from the surroundings into the reaction
system.
An exothermic process (q < 0) is one in which heat
flows out the reaction system into the
surroundings.
Energy Diagrams
• Most reactions need a “kick-start” to
occur, referred to as its “activation
energy”.
• The potential energy of a reaction
increases as the reaction progresses as
bonds are broken.
• At the point of maximum energy all
substances are broken apart in “activated
complex”.
• This complex is unstable and substances
recombine into lower energy products.
Reaction Rate Factors
• The rate of any reaction can be modified
by changing the conditions.
o collision theory can help explain why these
factors change the reaction rate
• Several strategies can be used to speed
up reactions:
o
o
o
o
Increase the temperature
Increase the concentration
Decrease the particle size
Employ a catalyst
Temperature
• Increasing the temperature speeds up the
reaction, while lowering the temperature
slows down the reaction
• Recall that temperature is directly proportional to
average kinetic energy
• At a higher temperature, there are more
effective collisions (more molecules have
the kinetic energy needed to react)
– For every 10oC increase in temperature, the
reaction rate usually DOUBLES!
Concentration
• Higher concentration of reacting particles
creates a faster rate of reaction.
• Increasing the concentration, increases
the frequency of the collisions, and
therefore increasing the reaction rate.
• ***For gases, increasing pressure and/or
decreasing volume increases concentration.
Particle Size
• The smaller the particle size, the larger
the surface area for a given mass of
particles. Effectively is increasing the
concentration.
• An increase in surface area increases the
amount of the reactant exposed for
collision to take place…
– Which increases the collision frequency
and the reaction rate.
• Methods:
– Grinding the reactants into a powder
– Dissolving in a solvent.
Catalyst
•
A catalyst is often the best way to speed
up an reaction.
•
In fact, some reactions simply will not go
forward measurably without one.
•
A catalyst is a substance that increases
the rate of a reaction without being
changed or used up during the reaction
•
The key is that they permit reactions to
proceed at lower activation energy than
is normally required
Nature of Reactants
• The type of reactant substances also plays
a role in reaction rates.
• Weak bonds = easier to break, reactants
with weak bonds will react faster.
– Essentially results in a low activation energy.
• Strong bonds = harder to break, reactants
with strong bonds will react slower (high Ea)
• Electronegativity and ionization energy also
plays a role.
Acids & Bases
• Properties and operational definitions of
acids and bases.
•
•
•
•
•
•
Kw = [H3O+][OH-]= 1x10-14
pH = -log [H3O+]
“pH loop”
Strong vs. weak acids and bases (Ka)
Salt hydrolysis
Arrhenius and Bronsted-Lowry models