Ludtke 1
Honors Chemistry Study Guide. Units 11-17
Leo Taffe, Vijay Ravindran, Jimmy Wiaduck, and Ben Ludtke
UNIT 11
Modern Atomic Theory
Key Terms:
Wavelength - The distance between two waves
Frequency - The number of waves that pass a given point per unit of time
Photon - Tiny packet of energy
Excited State - A state that contains an electron with excess energy
Ground State - The lowest possible energy state for an atom
Orbital - The region where an electron resides in an atom
Pauli Exclusion Principle - An atomic orbital can hold no more than two electrons,
and they must be of opposite spin
Valence Electrons - Electrons in outermost occupied principal energy level
Core Electrons - The inner electrons of an atom
Ionization Energy - The energy required to remove an electron from an atom in the
gas phase
Chapter Summary:
Rutherford's model of the atom
 Atoms are so small that they are difficult to study directly
 Atomic models are constructed to explain experimental data on collections
of atoms
 He discovered this through his golf foil experiment; there is a small,
positive, and dense region known as the nucleus. Most of the atom is empty
space (electrons)
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Electromagnetic Radiation
 Light has 2 natures: it can exist as a particle of energy (photon or quanta) or
as a wave
 Electromagnetic radiation is the whole spectrum of light, not just visible
light
 The different types of waves are: radio waves, microwaves, infrared, visible
light, ultraviolet, x-rays, and gamma rays
 The shorter the wavelength, the more powerful the wave is
 A wavelength is the length of one wave and is represented by "λ"
 Different waves travel at different speeds
 The frequency of a wave determines the length of the wave. Inversely
proportional
 Frequency is represented by "v"
 Speed = wavelength (in meters) x frequency per second
 C = λv
C = 2.998 x 10^8 m/s
 Max Planck discovered that energy and frequency are directly proportional
 E = hv
Planck's constant (h) = 6.626 x 10^-34 J x S
 Planck called the energy packets "quanta"
Atoms and Light/Energy
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Atoms can give off light.
They first must receive energy and become excited.
The energy is released in the form of a photon.
When an electron is excited and filled with energy, it goes to the excited
state
 When it releases its energy in the forms of photons of light, it falls back
down to the ground state
 Bohr’s model of the atom --> Quantized energy levels, electrons move in a
circular orbit, and electrons jump between levels by absorbing or emitting
photon of a particular wavelength
 Electrons don't move in a circular orbit, but they do reside in orbitals
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Orbitals and Energy Levels
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Orbitals don't have sharp boundaries
All atoms contain energy levels. There are 7 different energy levels
The more electrons an atom has, the more energy levels it has
The energy levels are divided into sublevels (s, p, d, and f)
2 electrons with opposite spins can fit in 1 orbital
The s sublevel only contains 1 orbital, the p sublevel contains 3 orbitals, the
d sublevel contains 5, and the f sublevel contains 7
 Each row on the periodic table has sublevels
Electron configuration
 The electron configuration determines elemental properties.
 It is the arrangement of electrons (for example, for Ne - 1s2, 2s2, 2p6)
 Noble gas shorthand - using a noble gas before the next set of levels to
decrease the amount of writing needed (Ex: At = [Xe] 6s2 4f14 5d10 6p5)
 ** always use the diagonal rule when determining the electron configuration
1s
first principal energy level
2s 2p
second principal energy level
3s 3p 3d
third principal energy level
4s 4p 4d 4f
fourth principal energy level
5s 5p 5d 5f
fifth principal energy level
6s 6p 6d
sixth principal energy level
7s
seventh principal energy level
 Orbital diagrams - use the amount of electrons and put them over the
sublevels in order. Use arrows to show the spins
 Core electrons are the inside electrons, valence electrons are the outside ones
and are typically the ones that react. Atoms with the same valence electron
arrangement often times behave in a similar manner
 Metals lose electrons and form cations, nonmetals gain e-s and form anions
 The size of an atom decrease across a row and increase down a column
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 The ionization energy, however, increases across a row and decreases down
a column
Important names and equations:
Ernest Rutherford
Max Planck
Niels Bohr
Albert Einstein
Schrodenger
c = λv
E = hv
*** NOTE: You can manipulate these equations to make a different equation that
will help you solve the problem
Ex: If c = λv and E = hv, then λ = ch/E
Review Questions:
1. As the wavelength increases, the frequency _____________
a) increases b) decreases c) doesn't change d) cannot be determined
2. When an electron in an atom absorbs energy, it _____________
a) moves to a higher energy level b) emits light c) moves to the ground state
3. Which of the following has the highest ionization energy?
a) K
b) Ge
c) Si
d) O
4. Which neutral atom has seven 3d electrons? _____________
5. Write the electron configurations for: Silicon, Selenium, and Oxygen. You may
use noble gas shorthand.
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UNIT 12
Chemical Bonding
Key Terms:
Bond - A force that holds atoms together
Bond Energy - The energy required to break a bond
Covalent Bond - Electrons are shared by 2 atoms
Ionic Compound - Formed by a metal and a nonmetal
Polar Covalent Bond - Electrons are shared unequally between 2 atoms
Electronegativity - The ability of an atom to attract shared electrons to itself
Resonance - More than one valid Lewis dot structure may be drawn for a molecule
Linear Structure - 2 electron pairs (2 bonds) on a central atom
Trigonal Planar Structure - 3 electron pairs (3 bonds) on a central atom
Tetrahedral Structure - 4 electron pairs (4 bonds) on a central atom
Trigonal Pyramid Structure - 3 single bonds and a lone pair on a central atom
Bent or V-Shaped Structure - 2 single bonds and a lone pair on a central atom
Chapter Summary:
Bonds
 The strong electrostatic forces of attraction holding atoms together in a unit
are called chemical bonds
 Ionic compound results when a metal reacts with a nonmetal
 A covalent bond results when electrons are shared by nuclei
 A polar covalent bond results when electrons are shared unequally by nuclei
 This is when one atom attracts the electrons more than the other atom
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 Electronegativity plays a large role in this. The more electronegative, the
more the pull on the electrons. It increases up and right
 If the difference in electronegativity between the 2 atoms is great, a polar
covalent bond can be formed
 If the difference is between 0.4 and 1.7, it is polar
 For example, HF is very polar
 Polar covalent molecules have dipole moments. This means that the center
for positive charge and the center for negative charge are separate
 Water has a dipole moment. This contributes to certain properties of water
Ions and Ionic Bonding
This is always helpful
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Ions are formed after metals or nonmetals transfer valence electrons
Atoms in stable compounds usually have a noble gas electron configuration
Metals lose electrons to get to this state, nonmetals gain electrons
Chemical compounds are always electrically neutral
Ions of opposite charges are packed together in a tight bond
Cations are smaller than the original atom, anions are larger
Polyatomic ions work the same as regular ions, it's just the atoms in a
polyatomic ion are bonded together covalently to act as a unit
Lewis Dot Structures
 ONLY include the valence electrons
 All atoms should be bonded so that they reach the octet rule (unless the
problem specifies otherwise)
 Bonding pairs are shared, lone pairs are not
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Single bond – covalent bond in which 1 pair of electrons is shared by 2
atoms
Double bond – covalent bond in which 2 pairs of electrons are shared by 2
atoms
Triple bond – covalent bond in which 3 pairs of electrons are shared by 2
atoms
A molecule shows resonance when more than one Lewis structure can be
drawn for the molecule (see below)
 There are very few exceptions to the octet rule (BF3)
Molecular Structure
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Three dimensional arrangement of the atoms in a molecule
Linear - Atoms in a line (ex: all molecules with only 2 atoms, CO2, etc)
Trigonal planar - Atoms in a triangle (BF3)
Tetrahedral - CH4
These last three are nonpolar molecules, meaning they are symmetrical
The VSEPR model (valence shell electron pair repulsion)
The molecular structure is determined by minimizing repulsions between
electron pairs
 Bent - Such as H20. It is NOT symmetrical. This is because there are lone
pairs present, and the lone pairs force the shared electrons down, causing
them not to be linear
 Trigonal Pyramidal - Such as NH3. There is one lone pair present
 **NOTE - The bond type does not matter when you are predicting the
overall shape of the molecule. A triple bond is counted the same as a double
or single bond
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Review Questions:
1. A bond formed between two elements that have a very large difference in
electronegativity is called _______________
a) A covalent bond b) A polar covalent bond c) A double bond d) An ionic bond
2. Which of the following molecules exhibits resonance?
a) NH3 b) SO3 c) CCl4 d) H2S
3. The shape of the NF3 molecule is ____________
a) Linear b) Bent c) Tetrahedral d) Trigonal Planar
4. Draw Lewis dot structures for each of the following molecules/ions and predict
the structure for each
a) NI3
b) CS2
c) COCl2 (C is the central atom)
d) NH4+
e) OF2
5. Draw Lewis structures for the carbonate ion, CO3-2, showing all possible
resonance structures:
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UNIT 13
Gases
Key Terms:
Absolute Zero - The temperature at which a gas theoretically occupies zero volume
Universal Gas Constant - 0.08206 Liters x Atmosphere / (mol x temperature)
Boyle's Law - The volume of a given amount of gas is inversely proportional to its
pressure
Charle's Law - The volume of a given gas is directly proportional to its temperature
in Kelvin
Avogadro's Law - Volume of a gas is directly proportional to the number of moles
of gas
Combined Gas Law - P1V1/T1 = P2V2/T2
OR
P1 V1 T2 = P2 V2 T1
Dalton's Law of Partial Pressures - The total pressure of a mixture of gases is equal
to the sum of each gas
Molar Volume of Ideal Gas Law - 22.4 L
Chapter Summary:
Gases
 Gases, unlike liquids and solids, are easily compressed because of the space
between the particles
 Under pressure, the particles in a gas are forced closer together.
 The amount of gas, volume, and temperature are factors that affect gas
pressure.
 Pressure, Volume, Temperature, and the number of moles are 4 very
important things in describing a gas
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 The more moles, higher pressure. Higher temperature, higher pressure.
Lower volume, higher pressure.
 Barometer – device that measures atmospheric pressure. It uses mm Hg
 Invented by Evangelista Torricelli in 1643
 1 standard atmosphere
= 1.000 atm
= 760.0 mm Hg
= 760.0 torr
= 101,325 Pa
 A manometer measures the pressure of a gas in a container.
 Boyle figured out that pressure and volume were inversely proportional
 Another way of stating Boyle’s Law is P1V1 = P2V2
 Charles found that volume and temperature were directly proportional
 His equation is V1/T1 = V2/T2 OR V1T2 = V2T1
 Avogadro's law --> V1/N1 = V2/N2
Ideal Gas Law
 This is the most crucial thing to know. It is the hardest to remember and
also the most important.
 Combining all of the other equations gives us this equation: PV = nRT
 In this equation, R is the universal gas constant (R = 0.08206)
 Volume in liters at STP = 22.4L
Partial Pressures
 the total pressure exerted is the sum of the partial pressures of the gases
present.
 Ptotal = P1 + P2 + P3
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The volume of the individual particles is not very important
The forces among the particles must not be very important
Total pressure is the pressure of the gas + the vapor pressure of the water.
Diffusion is the tendency of molecules to move toward areas of lower
concentration until the concentration is uniform throughout.
 During effusion, a gas escapes through a tiny hole in its container.
 Gases of lower mass move out more quickly than gases with higher masses
 Graham’s law of effusion states that the rate of effusion of a gas is
inversely proportional to the square root of the gas’s molar mass.
Ex: You have hydrogen gas (H2) and oxygen gas (O2). That's root of 32/4
 Study the Kinetic Molecular Theory for Ideal Gases
Reminders: ALWAYS convert any temperature given to you to degrees
Kelvin. Otherwise you'll fail. Use pressure in ATM and Volume in Liters or
else certain equations won't work. Concentration in moles, not grams or
anything else.
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Review Questions:
1. There are 11.2 Liters of an ideal gas at STP. The number of molecules of gas in
this sample is:
a) 3.01 x 10^23
b) 6.02 x 10^23
c) 11.2
d) 22.4
2. A sample of gas has a volume of 50.0 L at a temperature of 300K. What
temperature would be needed for this sample to have a volume of 60 L if its
pressure remains constant?
a) 400 K
b) 87 C c) 360 C d) 250 K
3. Under the same conditions of temperature and pressure, how many more times
faster will hydrogen effuse compared to ammonia?
4. 3.50g of CO2 and 11.6g of O2 are placed in a 3.00 L container at 273 K. What
will the total pressure in the container be?
5. CaCO3 decomposes when heated to form CaO and CO2. What volume of CO2
can be produced from 37.0g CaCO3 at a temperature of 25 C and a pressure of
1.15 atm?
6. A 54.0 mL sample of oxygen is collected over water at 23 C and 770 torr
pressure. What is the volume of the dry gas at STP? (Vapor pressure of Water at 23
C = 21.1 torr)
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Unit 14
Liquids and Solids
Key Terms:
Intermolecular Forces - Forces between molecules
Intramolecular Forces - Chemical bond between the atoms of molecules (bonds)
Dipole-Dipole Attraction - Attraction between polar molecules
Hydrogen Bonding - Strong dipole-dipole attraction including a hydrogen atom
London Dispersion Forces - Forces between nonpolar molecules
Molar Heat of Fusion - Energy required to melt 1 mole of a substance
Molar Heat of Vaporization - energy required to change 1 mole of a liquid to vapor
Evaporation - Transition from liquid to vapor phase
Condensation - Transition from vapor to liquid phase
Vapor Pressure - Pressure of a vapor in equilibrium with a sample of liquid
Crystalline Solid - Solid with a regular arrangement of components
Alloy - Matter that contains a mixture of elements and has metallic properties
Chapter Summary:
Intermolecular Forces
 Intermolecular forces are forces of attraction between particles
 Intermolecular forces are important in determining many macroscopic
properties of a substance
 Whereas intramolecular forces are inside of molecules, such as ionic
and covalent bonds, intermolecular forces occur on the outside
 There are four types of intermolecular forces. They are: ionic,
hydrogen, dipole-dipole, and London dispersion forces
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Ionic Bonding
 The strongest type of intermolecular forces
 Very high boiling point. NaCl is an example of this. Usually solid at room
temperature
Hydrogen Bonding
 Occurs between an H and a very electronegative atom (N, O, or F)
 Above is an example of a hydrogen bond between two water molecules
 Hydrogen bonding affects the boiling point, usually making it higher
 The boiling point increases because the forces are stronger, and therefore
take more energy to break
 That is why water is liquid at room temperature, and not a gas
Dipole-Dipole Attractions
 The force of attraction between opposite partial charge of neighboring
dipoles
 Hydrogen bonds are example of extremely strong dipole dipole attractions
London Dispersion Forces
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The force of attraction between induced dipoles
As the size of the atoms increase, the forces become stronger
For these forces, an electrically neutral atom can create its own dipole
Since I can't really explain it that well, here is a diagram
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Phase Changes
 Water and its phase changes; heating and cooling curve of water
 The density of water is 1g/mL. Ice happens to be less dense than water.
 Its boiling point is 100C and its freezing point is 0C
 It takes a certain amount of energy to vaporize a certain amount of water. It
also takes energy to melt it
Liquids
 Vaporization and Evaporation are not the same thing. Vaporization is a
process that occurs at the boiling point of a liquid, while evaporation can
occur at any given temperature.
 When the maximum amount of a liquid has evaporated at a given
temperature, the liquid has reached equilibrium and will remain constant
 Vapor pressure is the pressure of the vapor present at equilibrium with its
liquid
 At the boiling point, the atmospheric pressure is equal to the vapor pressure
Solids
 There are ionic solids, molecular solids, and atomic solids
 Ionic = Sodium Chloride (NaCl - Na+ + Cl-). These have high melting
points and are held together by strong forces between ions
 Molecular = Water (H2O). These melt at generally low temperatures. Held
together by weak intermolecular forces
 Atomic = diamond (composed only of C atoms). Properties vary greatly
 Metals are held together by closely packed atoms, also known as the electron
sea model. Alloys are metals that have a number of different atoms present
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Formulas and notes
*The molar heat of fusion and the molar heat of vaporization have formulas
*The molar heat of fusion for water is 6.02 kJ/mol; the molar heat of vaporization
is 40.6 kJ/mol (although you probably aren't expected to memorize either of these)
*Know the equation Q = sm t
*When you're using the equation, make sure to add all of the answers from the
equations. For example, if the questions asks you to find the amount of heat needed
to go from 50C to 150C, you must add the Q=sm t numerical answer to the
numerical answer for the molar heat of vaporization.
*As always, be able to convert from grams to moles and moles to grams for these
equations
Review Questions:
1. Which of the following substance most likely has the highest melting point?
a) NaCl b) H2O c) CH3OH d) Ne
2. Which of the following most likely has the lowest boiling point?
a) MgCl2 b) C12H22O11 c) H20 d) CS2
3. If the pressure above a liquid is increased, its boiling point will:
a) increase b) decrease c) stay the same
d) depends on its properties
4. Which substance has the highest vapor pressure?
a) H2O b) H2S c) CH4 d) KCl
5. How much heat is necessary to completely vaporize 75 g of ice at 0C? The
molar heat of fusion for water is 6.02 kJ/mol. The specific heat capacity of water is
4.18 J/g C. The molar heat of vaporization is 40.6 kJ/mol
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UNIT 15
Solutions
Key Terms:
Solution - Homogenous mixture in which the solute is dissolved in the solvent
Solvent - The major part of the homogenous mixture
Solute - The minor part of the homogenous mixture
Aqueous Solution - Homogenous mixture in which the solute is dissolved in water
acting as the solvent
Saturated - A type of solution that contains as much solute that will dissolve at a
given temperature
Unsaturated - A type of solution that has not reached the maximum amount of
solute able to dissolve at a given temperature
Supersaturated - A type of solution that became saturated at an elevated
temperature, then allowed to cool in a way that keeps all of the solid remained
dissolved
Concentrated - A situation where there is more solute than solvent able to
dissolve it in a 1:1 ratio
Dilute - A situation where there is more solvent than solute able to be dissolved in
a 1:1 ratio
Mass Percent - The percentage of mass of solute compared to mass of whole
solution
Molarity(M) - The amount of moles of a substance per liter(volume) of the
substance
Standard Solution - A solution whose concentration is accurately known in a lab
setting
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Dilution - A process in a lab setting to make the concentration of the solution
lower by the addition of water acting as a solvent
Neutralization Reaction - A reaction involving an acid and a base
Equivalent of an Acid - The amount of acid that furnishes(creates) 1 mole of H+
ions
Equivalent of a Base - The amount of base that furnishes(creates) 1 mole of OHions
Equivalent Weight - The mass in grams of one equivalent of an acid or base
respectively
Colligative Property - A property present in a solution that is based on the ratio of
the number of solute particles to the number of solvent particles, rather than the
type of particles present in the solution (ex: Boiling point elevation or freezing
point depression)
Chapter Summary:
Solution Basics
• Solution = Solute + Solvent
• Polar substances dissolve other polar substance, while non-polar
substances dissolve other non-polar substances. "Like dissolves Like"
• When a substance is dissolved, either intermolecular or intramolecular
forces are broken and new bonds are formed.
Solubility of Ionic Substances
• Ionic substances best dissolve in aqueous solutions
• Ionic substances break up into individual anions and cations when
dissolved due to the great difference in charges between the anions and
cations, combined with the attractive pull of the polar and non-polar sides
of the H2O molecule
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Ex: NaCl(s) + H20(l) --> (Na+) + (Cl-)
Solubility of Polar Substances
• Polar substances are molecules with a charge. H20 is the universal solvent
due to its properties of being an extremely polar molecule.
• When a polar solid dissolves in water, it often has to breaks apart from the
other intermolecular forces connecting it to the other molecules of the
polar solid.
Ex: Sugar molecules bond together through hydrogen bonding. The
Hydrogen bonds are broken during the dissolving process due to the
polar sides of the of the water molecule attracting different sides of the
polar molecule.
• Most alcohols are soluble in water, due to their polarity caused by the OH
bond.
Insolubility in Water
• Non-polar substances are insoluble in water. They are only soluble in other
polar substances.
Ex: Oils are not soluble in water because they are non-polar. They are
soluble in other non-polar substances such as paint thinner.
Rate of Dissolving
• Stirring a substance that is dissolving increases the rate of dissolving due to
the increased speed and likelihood of molecules hitting each other.
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• Grinding a solute before putting it in a solvent to be dissolved, increases the
solids surface area, which therefore increases the speed of the dissolving.
This is because the increased surface area allows for more chemical
processes to occur at once.
• Heating during the process of dissolving a substance increases the rate of
dissolution due to the fact that heated molecules move faster, which
therefore allows for more chemical processes to occur. Vice versa is true if
cooling takes place during the dissolving process.
General Facts
• Dissolving a solute in a solution will decrease the vapor pressure of the
solution due to the lack polar molecules on the surface
• Adding salt to water increases the boiling point of the water
• Saturated solutions can be concentrated or dillute
Math Concepts
• Molarity = (moles/Liters)
• Mass Percent = (mass of solute/mass of solution) x 100%
= [mass of solute/(mass of solute) + (mass of solvent)]
x 100%
• Moles of Solute = (M1)(V1) = (M2)(V2)
(M1) = Molarity of initial solution (V1) = Volume of initial solution
(M2) = Molarity of diluted solution (V2) = Volume of diluted solution
Example Problems
1.) 70.0 mL of .15M Ag(NO3) solution is poured into 140.0 mL of 0.10M MgCl2
solution. What mass of solid AgCl will be formed?
2.) 22.5g of MgCl2 is dissolved in enough water to make 500. mL of solution. The
concentration of Cl(-) ions is what?
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UNIT 16
Acids and Bases
Key Terms:
Acids - A molecule with a hydrogen ion
Bases - A molecule with a hydroxide ion
Arrhenius Concept of Acids and Bases - Produces either an H+ or OH- for an acid
or base respectively
Bronsted-Lowry Model - Acids and bases are proton donors and acceptors
Conjugate Acid - Molecule with one more hydrogen ion compared located on the
product side of equation compared to the acid on the reactant side
Conjugate Base - Molecule with one less hydrogen ion on the product side of the
equation compared to the base on the reactant side
Hydronium Ion - H30+
Completely Ionized - All parts of the products completely disassociated
Strong Acid -HCl, HBr, HI, H2(SO4), H(NO3), HClO4 (Completely ionized acids)
Strong Base - Base that completely dissociates
Dipotic Acid -Acid containing two hydrogen ions
Oxyacids -Acids with proton attached to an oxygen atom
Organic Acids -Hydrogen ion attached to a molecule with Carbon
Carboxyl Group -Group of Carbon, Oxygen atoms, and Hydrogen ion found in
organic acids
Amphoteric Substance - Substance that can be an acid or base (water)
Ion-Product Constant- [H+][OH-] = 1 x 10^-14
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pH Scale - Scale used to indicate the acidity of a solution
Indicators - Things that indicate the acidity or how basic a solution based on color
Indicator Paper - Paper that changes color based on pH
Neutralization Reactant - The reactants of a reaction involving an acid and a base
Titration - Delivering a measured volume of a solution with a known
concentration, into the solution being analyzed
Buret - Device used for accurate measurement of the delivery of a liquid
Stoichiometric Point (Equivalence Point) - When exactly enough standard solution
(titrant) has been added to react with all the solution being analyzed
Titration Curve (pH Curve) - Data plotted on a graph for a given titration showing
the pH of the solution vs volume
Buffered Solution - Resists a change in its pH when either an acid or base has
been added. Presence of a weak acid and its conjugate base or weak base and
conjugate acid can act as a buffer.
Neutralization Reaction
• HCl is the acid in this equation and H20 is acting as a base.
• H30(+), also known as hydronium, is acting as the conjugate acid in
this equation. Cl- is acting as the conjugate base.
• As shown above, HCl and Cl- form an Acid-Conjugate base pair, while
H20 and H30+ form a Base-Conjugate acid pair
• In general, if a weak acid is on the reactant side of the equation, they
will have a strong acid for a conjugate acid. If a strong acid is on the
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reactant side of the equation, its conjugate acid will be weak. These
same rules can also be applied to bases and conjugate bases. That is
why in the diagram above the strong acid HCl has a weak conjugate
acid.
• The forward reaction dominates the example above because the
reactant side contains a strong acid being HCl. If the reactant side
contained a weak acid instead of a strong acid, the reverse reaction
would dominate.
Solutions
• [H+] = Hydrogen molecule concentration, [OH-] = Hydroxide molecule
concentration, Kw = 1 x 10^-14 = equilibrium constant for concentration of
water
• [H+] > [OH-] is an acidic solution -[H+] < [OH-] is a basic solution - [H+] =
[OH-] is a neutral solution
• [H+][OH-] = 1 x 10^-14 = Kw at 25 degrees Celsius
• pH = -log[H+]
pOH = -log[OH-]
• pH + pOH = 14.00
Titrations
• M(a) = Molarity of Acid
V(a) = Volume of Acid
• M(b) = Molarity of Base
V(b) = Volume of Base
• [M(a)][V(a)] = [M(b)][V(b)]
Example Problems
1.) If 60.0 mL of a HBr solution requires 35.0 mL of a 0.30 M NaOH to titrate it to
the equivalence point, what is the concentration, in moles per liter, of the HBr
solution?
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2.) What is the pH of a 0.15 M H(PO4) solution?
UNIT 17
Equilibrium
Key Terms:
Collision Model - The colliding of molecules is what is needed in order for a
reaction to occur
Activation Energy - Energy required to be used in order for a reaction to start
Catalyst - A substance that speeds up a reaction without being consumed
Enzyme - Catalyst in a biological system
Homogenous Reaction - All reactants and products are in one phase
Heterogenous Reaction - Reactants are in two or more phases
Equilibrium - The exact balancing of two processes, one which is the opposite of
the other
Chemical Equilibrium - A dynamic state where concentrations of all reactants and
products remain constant. The rates of the reverse and forward reactions are
exactly the same.
Law of Chemical Equilibrium - A dynamic reversible state in which rates of
opposing processes(forward and reverse reactions) are equal
Equilibrium Expression - K = (Products) / (Reactants) = [(C)^c x (D)^d] / [(A)^a x
(B)^b]
Equilibrium Constant - The set ratio of products to constants, that stays this ratio
at equilibrium
Equilibrium Position - Idea of a theoretical position in the system based off the
value of the equilibrium concentration
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Homogenous Equilibria - An equilibria system where the products and reactants
are in the same state
Heterogeneous Equilibria - The products and reactants aren't in the same state in
an equilibria system
Le Chatelier's Principle - When a change is imposed on a system at equilibrium,
the position of the equilibrium shifts in a direction that tends to reduce the effect
of that change
Solubility Product Constant (Ksp) - Measures the solubility of a solution using the
equation Ksp = [M+][X-] with M and X as solids in a reversible reaction
Collision Model
• In order for a collision to successfully break a bond, and therefore trigger a
reaction, there needs to be enough energy to break the bond, the
molecule must collide, and wen the molecule collide they must be aligned
properly.
• Due to the laws of the collision model, the rate of the chemical reaction is
determined by the temperature of the system and the concentration of
molecules within the system.
• The concentration of a system is determined by the pressure exerted on it
and the volume it is allowed to encompass. Pressure and volume are
inversely proportional, because as the volume gets bigger, the amount of
pressure in the system goes down. The more pressure that is exerted on a
system, the faster the rate of the reaction becomes.
• The temperature of a system also affects the rate of a reaction by directly
affecting how fast molecules move. The higher the temperature the faster
the molecules move, the faster the molecules move and also the more
often molecules bump into each other. This results in more positive
reactions between molecules to take place, and overall a faster reaction of
the system. If the temperature of a system is cooler than normal, vice
versa occurs in regards to the effects and rate of the overall reaction.
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• Catalysts lower the activation energy needed for a positive reaction to
occur.
Equilibrium Reactions
• They have the ability to move in both the forward and reverse direction.
When both of the forward and reverse reactions are occurring at the same
rate, equilibrium occurs.
• In accordance with Le Chatlier's principle, when a change is exerted on the
system, the system shifts in the direction that will most effectively alleviate
the effects of the change and bring back the system to equilibrium.
• The constant(K), resulting from an equilibrium expression of reaction, must
always remain constant. It can do so by the equation shifting to balance
the reactants and the products. The number of moles of each of the
products and reactants may change, but the concentration will always stay
the same in order for the reaction to remain at equilibrium. Solids and
liquids aren't included in the expression, only gases and aqueous solutions.
• Adding more of a reactant to a system will cause there to be more
products and a little bit less of the other reactants. Therefore, the reaction
shifts right.
• Increasing the volume of a system will cause the reaction to shift towards
the side where there is more moles of gas/aqueous solutions.
• Decreasing the volume of a system will cause the reaction to shift towards
where there is less moles of gas/aqueous solutions.
• Decreasing the pressure causes the reaction to shift towards where there
is more moles of gas/aqueous solutions.
• Increasing the pressure causes the reaction to shift to where there is less
moles of gas/aqueous solutions.
• In an equilibrium system, it always shifts away from where heat is added.
In an exothermic equation, the temperature is decreased or created so
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heat shows up on the product(right) side. In an endothermic reaction, heat
is used up in the reaction, so it shows up on the reactant(left) side of the
reaction.
• If Heat = -�H then the reaction is exothermic and shifts left.
• If Heat = +�H then the reaction is endothermic and shifts right.
• If the temperature is "increased", then the reaction shifts left and is
exothermic.
• If the temperature is "decreased", then the reaction shifts right and is
endothermic.
Equilibrium Expressions
• K = Equilibrium Constant
• K never changes
• K > 1 means equilibrium position is far to the right
• K < 1 means equilibrium position is far to the left
• K = (Products) / (Reactants) = [(C)^c x (D)^d] / [(A)^a x (B)^b]
• Ksp = Solubility Product Constant
• Ksp = [M+][X-] with M and X as solids in an equilibrium expression
• If a coefficient is in front of one of the solid ions, it is incorporated in to the
solubility product constant as an exponent.
Example Problems
1.) Using the equation 5A(s) + 7B(g) <--> C(aq) + 4D(g) calculate equilibrium
constant if [A] = .0025 M, [B] = .450 M, [C] = 1.2 M, and [D] = 0.30 M.
2.) Ksp for MgF2 = (3.7)10^-8 at 25 degrees Celsius. Calculate the solubility of
MgF2 at this temperature.
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EXAMPLE PROBLEMS: THE WRITTEN RESPONSE QUESTIONS
**Note - This portion is nearly 40% of the test. Each of these 5 question is
10x more important than each multiple choice. Here is an example of each.
1. (Electromagnetic Radiation Calculation) A wave of visible light has a
wavelength of 5.93 x10^-7m. What is the energy of a single photon of this
wave?
2. (Gas Law Calculation) Find the number of grams of CH4 that exert a
pressure of 1.80 atm at a volume of 32.5L and a temperature of 50.0C?
3. (Titration Calculation) If 60 mL of a HCl solution requires 22.0 mL of
0.30 M NaOH to titrate it to the equivalence point, what is the concentration,
in molarity, of the HCl solution?
4. (Solution Stoichiometry) 40.0 mL of 0.20 M AgNO3 solution is poured
into 200.0 mL of 0.08 M MgCl2 solution. What mass of solid AgCl will be
formed?
5. (Balance Redox Reaction) Using the half reaction method, balance the
following reaction.
CeO2 + Sn+2 + H+ -----> Ce+3 + Sn+4 + 2H20