Study Guide for Honors Chemistry Final - June 2015
You will be given all needed equations, values of constants and other pertinent information, including the
following: periodic table, c = , E=h, the value for h = 6.626x10-34 J.s, the conversion between meters and
nanometers (1 m = 1 x 109 nm), electronegativity values, activity series of metals, gas law equations, dilution
equation, molarity, molality. It is your responsibility to recognize the equations and understand how to use them.
You will also be given a 3” x 5” notecard that, must be handwritten by yourself, you may use to organize your
thoughts for the exam.
Do the sample problems and review all the vocabulary words in the sections indicated.
Go to the following links to test your knowledge on the Chapters (remember we did not do all of the
sections in all of the following chapters):
http://chemistrymc.com/chapter_test and http://chemistrymc.com/standardized_test
Electrons in Atoms – Chapter 5
Objectives:
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Describe the relationship between wavelength and frequency.
Be able to use the equation c =  where c = 3.00x108 m/s
Describe the relationship between frequency and energy of a wave.
Be able to use the equation E = h where h = Planck’s constant = 6.626x10-34 J.s
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Explain the difference between a continuous spectrum and a line spectrum and give examples of
each.
Be able to arrange the colors of the visible spectrum from lowest energy to highest energy.
Define a quantum of energy and explain how it is related to an energy change of matter.
Describe how the quantum mechanical model describes what an atom is and where the subatomic
particles are located.
Explain the relationships among an atom’s energy level, sublevels, and atomic orbitals.
Apply the Pauli Exclusion Principle, the Aufbau Principle, and Hund’s Rule to write electron
configurations using orbital diagrams and electron configuration notation.
Define valence electrons and draw electron-dot structures representing an atom’s valence electrons.
Be able to read and write quantum numbers given an electron configuration or an element’s symbol
Labs – Thickness of Zinc, Spectroscopy and Flame Tests, Electron Arrangements
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Describe the contributions of Bohr, Schrödinger, Planck, Einstein, & deBroglie to determining the location of
electrons in the atom and their properties
Vocabulary:
 Wavelength, frequency, amplitude and speed
of waves
 Electromagnetic spectrum
 Quantum
 Photoelectric effect
 Atomic emission spectrum
 Ground state vs. excited state
 Heisenberg uncertainty principle
 Quantum mechanical model of the atom
 Atomic orbital
 Principal quantum number
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Principal energy level
Orbitals (s, p, d, and f)
Electron configuration
Valence electron
Electron-dot structures (Lewis dot diagrams)
Quantum numbers
n, l, ml, ms
Periodic Table and Periodic Trends – Chapter 6
Objectives:
 Identify key features of the periodic table, representative elements, alkali metals, alkaline earth metals,
halogens, noble gases, transition metals, inner transition metals, metals, nonmetals, metalloids/semimetals,
groups vs. periods, formation of ions.
 Know what the contributions of Mendeleev and Moseley are to the development of the periodic table.
 Explain why elements in the same group have similar properties.
 Lab – Periodic Trends
Vocabulary:
 Periodic law
 Octet rule
 Group and period
 Ion
 Alkali metal, alkaline earth metal, transition
 Ionization energy
element (metal), inner transition metal,
 Electronegativity
halogen, noble gas, representative element
 Metalloid (semimetal), nonmetal, metal
 s-block, p-block, d-block, f-block
Chemical Names and Formulas – Chapter 8 – Section 8.3 & Chapter 9 – Section 9.2
Objectives:
 Distinguish between ionic and molecular (covalent) compounds.
 Write formulas for ionic compounds and oxyanions.
 Name ionic compounds and oxyanions.
 Know the names and formulas for the common polyatomic ions.
 Know the elements that exist as diatomic molecules.
 Be able to write the names and formulas of given compounds.
 Use the periodic table to determine the charge on an ion.
 Apply the rules for naming and writing formulas for ionic and molecular compounds.
 De able to define or describe: ion, cation, anion.
 Lab – Conductivity of Solutions
Chapter 9 – Covalent Bonding
Objectives:
 Use electron dot diagrams to show the structure of simple covalent compounds and polyatomic ions.
 Describe and give examples of resonance structures.
 Give examples of three classes of exceptions to the octet rule.
 Use VSEPR theory to predict the shape of simple covalent molecules.
 Describe what hybrid orbitals are and how hybridization relates to the shape of covalent molecules.
 Use electronegativity values to classify a bond as nonpolar covalent, polar covalent, or ionic.
 Determine whether a chemical bond is polar or nonpolar, AND whether a molecule is polar or
nonpolar.
 Lab – Shape and Polarity of Covalent Molecules – Determine the shape and polarity of molecules
using Lewis dot structures, VSEPR theory, and bond polarity.
Vocabulary:
 single, double and triple covalent bonds
 trigonal pyramidal
 structural formulas
 tetrahedral
 coordinate covalent bonds
 dipole
 VSEPR theory
 polar, nonpolar
 linear
 electronegativity
 trigonal planar
 sigma bonds
 trigonal bipyramidal
 pi bonds
 octahedral
 bent
Chapter 10 – Chemical Reactions (Predicting products and writing complete, ionic and net ionic
equations) and Chapter 20 – Redox Reactions and 16.2 – 16.4 – Heat in Chemical Reactions and
Hess’s Law
Objectives:
 Write equations describing chemical reactions using appropriate symbols.
 Write balanced chemical equations when given the names or formulas of reactants and products.
 Identify a reaction as combination (synthesis), decomposition, single-replacement, doublereplacement neutralization or combustion.
 Determine the products (if any) of single-replacement reactions using an activity series.
 Identify three kinds of products of double-replacement reactions.
 Know the products of complete and incomplete combustion of hydrocarbons.
 Write and balance complete, ionic and net ionic equations of reactions in aqueous solutions and
identify spectator ions.
 Describe the processes of oxidation and reduction.
 Determine the oxidation number of an element in a compound.
 Determine what was oxidized and what was reduced in a redox reaction.
 Explain the meaning of enthalpy and enthalpy of reaction and what a positive or negative sign means
 Use Hess’s law of summation of enthalpies of reaction to calculate the enthalpy change for a
reaction
 Lab – Types of Chemical Reactions – Be able to identify the types of chemical reactions based on
reactants AND be able to predict the products of the different types of chemical reactions.
Vocabulary:
 Chemical equation
 Catalyst
 Coefficients
 Balanced equation
 Combination reaction
 Decomposition reaction
 Single-replacement reaction
 Double-replacement reaction
 Combustion reaction
 Neutralization reaction
 Activity series
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Complete ionic equation
Spectator ions
Net ionic equation
Redox reaction
Oxidation
Reduction
Oxidizing agent
Reducing agent
Enthalpy, Enthalpy of Reaction
Hess’s Law
Chapter 13 – States of Matter
Objectives:
 Relate kinetic-molecular theory to some common properties of gases.
 Calculate the partial pressure of a gas, given appropriate information (Dalton’s Law).
 Describe how mass affects the rates of diffusion and effusion of gases (Graham’s Law).
 Describe and compare the behavior of molecules in the solid, liquid and gaseous state.
 Convert pressure measurements between mmHg, atm and kPa.
 Describe and compare types of intermolecular forces.
 Describe how hydrogen bonding explains some unique properties of water.
 Explain what is meant by the normal boiling point of a liquid.
 Vapor pressure diagrams, phase diagrams and cooling/heating curves of substances.
 Labs – States of Matter Lab, Liquid Nitrogen
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Vocabulary:
 Kinetic energy
 Kinetic-molecular theory
 Elastic collision
 Gas pressure
 Vapor pressure (of a liquid)
 Hydrogen bond
 Dipole-dipole attraction
 London Dispersion forces
 Atmospheric pressure
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Evaporation, Boiling, Vaporization,
Condensation
Melting, Freezing
Melting Point and Freezing Point
Boiling Point
Sublimation, Deposition
Crystalline solid
Metallic solid
Phase diagram
Triple point
Chapter 14 – Gases (14.4 covered with Stoichiometry) & Section 13.1
Objectives:
 Explain how the kinetic energy of gas particles relates to Kelvin temperature.
 Describe the relation between volume, pressure, temperature and number of moles of a gas sample.
 Apply the combined gas law to problems involving the temperature, volume, and pressure of a
contained gas.
 Use the ideal gas law in calculations involving the temperature, volume, pressure and number of
moles of gas.
 Lab – Molar Volume of a Gas
Vocabulary:
 Effusion
 Diffusion
 Graham’s Law of Effusion
 Pressure
 Dalton’s Law of Partial Pressures
 Barometer
 Pascal
 Atmosphere
 Ideal Gas
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Boyle’s Law
Charles’s Law
Gay-Lussac’s Law
Graham’s Law of Effusion
Combined Gas Law
Molar Volume
Avogadro’s Principle
Ideal Gas Law
Chapter 15 – Solutions
Objectives:
 Define concentration of a solution in terms of molarity.
 Describe the relationship of the polarity of solutes and solvents to solubility.
 Describe the effect of temperature on the solubility of solids and gases in water.
 Calculate the concentration of solutions.
 Calculate volumes and concentrations involved in the dilution of a solution.
Vocabulary:
 Solute
 Solvent
 Immiscible
 Miscible
 Solvation vs. Hydration
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Saturated solution
Solubility
Concentration
Dilute
Unsaturated solution
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Supersaturated solution
Concentrated
Molarity (M)
Percent by mass
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Percent by volume
Colligative property
Chapter 19 – Acids & Bases
Objectives:
 Identify the physical and chemical properties of acids and bases.
 Classify solutions as acidic, basic, or neutral.
 Compare and contrast the Arrhenius, Bronsted-Lowry, and Lewis models of acids and bases.
 Relate the strength of an acid or base to its degree of ionization.
 Explain the relationship between the strengths of acids and bases and the values of their ionization
constants.
 Write chemical equation for neutralization reactions.
 Explain how neutralization reactions are used in acid-base titrations.
 Activity – Titration – Determine the concentration of an acid or base using a titration.
Vocabulary:
 Acidic solution
 Basic solution
 Arrhenius acids & bases
 Bronsted-Lowry acids & bases
 Lewis acids & bases
 Conjugate acid
 Conjugate base
 Conjugate acid-base pair
 Amphoteric
 Strong acid or base
 Weak acid or base
 Acid ionization constant
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Base ionization constant
Ion product constant for water
pH
pOH
Neutralization reaction
Salt
Titration
Equivalence point
Acid-base indicator
End point
Laboratory Section:
These questions will cover the topics that we investigated in the labs we performed this second semester. It
will include questions about types of equipment that you used and safety practices that were covered during
the course of the year. Review the post-lab questions that were asked on the labs and understand the
purpose of each of the labs as well as what equipment that was used to make measurements.
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NAME: ____________________________________________
DATE: ___________________
PRACTICE PROBLEMS BY CHAPTER – Chemistry L1
Chapter 5 - Sample Problems:
1. What is the frequency of green light, which has a wavelength of 4.90 x 10-7 m?
2. An x-ray has a wavelength of 1.15 x 10-10 m. What is its frequency?
3. Which type of electromagnetic radiation has the most energy: Microwaves, X-Rays, or Visible
light? Why?
4. What is the energy of a type of radiation with a frequency of 9.50 x 1013 Hz?
5. What element is the name of the element that has the electron configuration of [Kr]5s24d105p1?
6. What element has the following electron configuration: 1s22s22p63s23p64s23d104p65s24d105p66s2?
7. Write out the electron configurations for the following elements: (For the exam make sure you
know how to do both the complete and the noble gas short cut for electron configurations)
a. Tc
b. Aluminum
8. Why are Cr and Cu exceptions to writing the electron configurations of elements?
9. For the following elements; state how many valence electrons they have and draw their electron-dot
symbol (Lewis dot symbol)
a. Ca
b. As
c. Sn
10. Using the following quantum numbers, what is the maximum number of electrons that the atom
could have?
n=3, l = 1 ml = -1, 0, 1
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Chapter 6 - Sample Problems:
1. Compare & contrast how Mendeleev’s periodic table was organized to how Moseley organized it.
2. What is the periodic law?
3. Identify each of the following as a metal, nonmetal, or metalloid and as representative elements or
transition elements:
a. Neon
b. Arsenic
c. Zinc
d. Magnesium
4. Match each numbered item on the right with the lettered item that it is related to on the left:
a. Alkali metals
1. group 18
b. Noble gases
2. group 17
c. Alkaline earth metals
3. group 2
d. Halogens
4. group 1
5.
Rank the following elements by increasing atomic radius: carbon, aluminum, oxygen, potassium
6.
Why does fluorine have a higher ionization energy than iodine?
7. Why do elements in the same family generally have similar properties?
Chapters 8 & 9 Chemical Names and Formulas Practice Problems:
1. Write the symbol and name for the ion formed when:
Ex: A potassium atom becomes an ion
__K+___ __potassium ion___
a. An aluminum atom becomes an ion
_______ ______________________
b. A fluorine atom becomes an ion
_______ ______________________
c. A nitrogen atom becomes an ion
_______ ______________________
2. Write the formula (including charge) for each polyatomic ion:
a. Hydroxide __________
b. Nitrate __________
c. Sulfate __________
3. Write the formulas for the following compounds:
a. Potassium sulfate _______________
d. Magnesium hydroxide _______________
b. Carbon tetrabromide _____________
e. Diphosphorus trioxide ________________
c. Silver chloride _______________
f. Iron(II) oxide ____________________
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4. Name the following compounds:
a. CS2
_____________________
d. Cu2O __________________________
b. NaI _______________________
e. K2S ___________________________
c. NH4Br _____________________
f. N2O5 ___________________________
Chapter 9 – Covalent Bonding Sample Problems:
1. Draw electron dot structure and state the shape of each of the following:
a. PCl3
c. BF3
b. CO2
d. CO32-
2. State the type of bonds (nonpolar covalent, polar covalent, or ionic) found in the following
compounds, and circle the polar molecules.
a. CO
c. SF6
b. NH3
d. CaO
Chapter 10 Chemical Reactions Sample problems:
Balance the following & state what type of reaction it is:
1. _____AgNO3 + _____H2S  _____Ag2S +_____HNO3
2. _____HCOOH + _____O2  _____CO2 + _____ H2O
3. _____Mg + _____ Zn(NO3)2  _____Mg(NO3)2 + _____Zn
4. _____SO2 + ______O2  _____SO3
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5. Predict the products in the following chemical reactions; make sure to balance the reactions.
a. _____ Zn (s) + _____ HCl (aq) 
b. _____ H2SO4 (aq) + _____ KOH (aq) 
6. Write a balanced complete ionic equations, a balanced net ionic equation for the following reaction
and state what the spectator ions are:
_____ Pb(NO3)2 (aq) + _____ H2SO4 (aq)  _____ PbSO4 (s) + _____ HNO3 (aq)
Complete ionic equation:
Net ionic equation:
Spectator ions:
Identify what is oxidized (O), what is reduced (R), and the oxidation numbers on all the elements.
1.
2Zn + O2  2ZnO
2. Mg + I2  MgI2
3. H2S + Cl2  S + 2HCl
Chapter 13 – States of Matter Sample Problems:
1. At 25°C, the vapor pressure of compound A is 40 mmHg and the vapor pressure of compound B is
50 mmHg. Which compound will evaporate faster? Which compound has the higher boiling point?
Explain.
2. Name two changes of state (e.g. freezing, melting, etc.) that require energy.
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3. State the three kinds of intermolecular attractive forces. Rank these forces from weakest to strongest
(for molecules of about the same size).
4. Which compound would you expect to have the higher boiling point, CH4 or NH3? Explain.
Chapter 14 – Gases and Section 13.1 (14.4 covered with Stoichiometry) Sample Problems:
1. Which gas diffuses faster, nitrogen or carbon monoxide? Why?
2. A gas at 155 kPa and 25oC occupies a container with an initial volume of 1.00 L. By changing the
volume, the pressure of the gas increases to 605 kPa as the temperature is raised to 125oC. What is
the new volume?
3. A sample of air has a volume of 6.00 L at 100 kPa. What volume will it occupy at 25.0 kPa if the
temperature remains constant at 30oC?
4. A steel cylinder that has a volume of 20.0 L is filled with nitrogen gas to a final pressure of 25.0 atm
at 28oC. How many moles of nitrogen gas does the cylinder contain?
5. A gas mixture containing oxygen, nitrogen, and carbon dioxide has a total pressure of 32.9 kPa. If
the pressure of oxygen is 6.6 kPa and the pressure of nitrogen is 23.0 kPa, what is the pressure of
CO2 in the mixture?
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Chapter 15 – Solutions Sample Problems:
1. 36.0 g of glucose are dissolved to make 2.00 liters of solution. What is the molarity of the solution?
(The molar mass of glucose is 180 g/mol).
2. How many grams of solute are in 335 mL of 0.425 M KNO3?
3. If 250 mL of 2.0 M CaCl2 is diluted to a final volume of 700 mL, what is the concentration of the
diluted solution?
4. Calculate the percent by mass of 3.55 g of NaCl dissolved in 88 g water.
5. What mass of water must be added to 255.0g KCl to make a 15.00 percent by mass aqueous
solution?
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Chapter 19 – Acids and Bases Sample Problems:
1. An aqueous solution tastes bitter and turns litmus blue. Is the solution acidic or basic?
2. Classify the following as an Arrhenius acid or an Arrhenius base:
a. H2S
c. Mg(OH)2
b. RbOH
d. H3PO4
3. Given the concentration of either hydrogen ion or hydroxide ion, calculate the pH, pOH and
concentration of the other ion at 298 K:
a. [H+] = 1.0 x 10-4 M
b. [OH-] = 1.3 x 10-2 M
4. Calculate [H+] and [OH-] in each of the following solutions at 298 K.
a. pH = 3.00
b. pH = 5.24
5. How many milliliters of 0.225 M HCl would be required to titrate 6.00 g KOH?
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