Study Guide for CHM 151LL Final – Written and Practical
Final Exam information:
The lab final will consist of multiple-choice questions and a practical portion.
• multiple choice (15 questions x 2 points each = 30 points)
• practical (13 questions = 70 points)
Combined counts as 20% of the final course grade
The lab final will begin with the multiple-choice portion. All materials will be provided for the final exam (test,
scan-tron, periodic table, scratch paper, calculator, and pencil – if needed). You will be allowed to use your lab
notebook but not any other papers. You will not be allowed to use a programmable calculator.
The practical will consist of several stations set up around the lab that will ask you to measure, complete short
experimental procedures, or observe an object or reaction and record results of what you see. There are also
sort answer questions with calculations to do at your desk.
The outline below is only intended to serve as a guide. The bullet points for each lab focus on the highlights of
procedures and calculations but do not go into detail. You should read through each lab procedure and re-work
all calculations from this past semester to fully prepare for this final.
Lab Safety:
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Review the safety label information (the four-color label that accompanies chemicals). What does each
color represent? What do the numbers in the diamonds represent?
Review the location of safety equipment in lab and the appropriate use of each.
Be familiar with safety procedures in lab (appropriate attire, waste disposal, cleanliness, contamination
of reagents, chemical spills clean-up, using Bunsen burners, etc.).
Density:
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Differentiate between accuracy versus precision. Accuracy is the difference between your measurements
and the accepted or true value; precision is the range of your measured values. Understand how to
calculate and evaluate both.
Determine which piece of glassware is more accurate and/or precise given data collected in lab.
Be able to read volumes in various pieces of glassware (buret, beaker, 100 mL graduated cylinder, and
10 mL graduated cylinder). How many significant figures can volumes for each piece of glassware
have?
Be able to calculate volume of a solid object using calipers.
Be able to measure the mass of an object on an analytical balance (by taring the weighing container or
subtracting the mass of the weighing container).
Calculate the density of an object by measuring mass and volume.
States of Matter:
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Identify and label pictures as atoms, molecules, compounds, pure substances, mixtures, solids, liquids,
and/or gases.
Be able to describe the position and motion of atoms in the solid, liquid, or gas state.
Understand, describe, and/or draw the difference between physical and chemical changes.
Given a scenario comparing a substance’s appearance before and after a change, determine if the
change was physical or chemical.
Atomic Theory:
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Understand what wavelengths represent on the Spec-20.
Be able to describe what each component in a Spec-20 does and how absorbance is measured.
Understand the relationship between wavelength, frequency, and the speed of light. Be able to perform
calculations using these variables.
Understand the relationship between frequency and energy and between wavelength and energy. Be able
to perform calculations using these variables.
How can substances be identified using flame tests? Why do different metals give off different colored
flames?
How are emission spectra obtained for elements? Describe the process electrons go through to produce
an emission spectrum.
Introduction to Spectroscopy: Analysis of Copper Ore:
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Calculate the concentration of a diluted solution: M1V1 = M2V2.
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Choose the appropriate wavelength to use on a Spec-20 given a graph of wavelength versus absorbance.
Why was it necessary to calibrate the Spec-20 between readings in Part II of the procedure but not in
Part III?
Be able to calculate the concentration of an unknown solution given a Beer’s Law Plot and or linear
equation.
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Molecular Geometry:
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Draw Lewis Structures (or electron-dot formulas) for a given compound by counting the valence
electrons for the atoms in that compound.
Determine the molecular geometry for a given structure or model.
Be able to identify the number of lone pairs of electrons on a central atom in a structure.
Determine if a formula needs resonance structures to accurately describe it.
Be able to identify when a structure violates the octet rule.
Identify the bond angles in a structure.
Use electronegativity differences between atoms to draw dipoles and determine the overall polarity of a
molecule.
Assign hybridization to the central atom in a molecule.
Intermolecular Forces:
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Given the formula or drawn structure of a compound, determine the primary type of intermolecular force
attracting two molecules together (London dispersion, dipole-dipole, hydrogen bonding).
Place substances in order of increasing strengths of intermolecular forces.
Place substances in order of increasing boiling points based on strengths of intermolecular forces.
Differentiate between bonds and intermolecular forces given a description or drawing.
Relate relative temperature differences during evaporation or rates of evaporation to strengths of
intermolecular forces.
Chemical Reactions:
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Predict products and typical observations of each type of reaction:
o Combination (one product is formed)
o Decomposition (heat is generally required, single reactant breaks down)
o Single-replacement (element + compound = element + compound)
o Precipitation (form a solid product)
o Acid-base neutralization (form water and salt; usually give off heat)
o Combustion (hydrocarbon reacts with oxygen gas to produce steam and carbon dioxide)
Predict phases of products made in reactions.
Balance equations.
Given an equation, determine the type of reaction that will occur.
Given the reactants of an equation, predict what type of reaction will occur, what the products will be,
what phases (s, l, g, aq) products will be in, and what the balanced equation will be.
Review nomenclature.
Know how to test solutions using litmus paper.
Predict products (including phases) of precipitation reactions.
Predict the formulas of precipitates formed in precipitation reactions.
Be able to write molecular, ionic, and net ionic equations for precipitation reactions.
Be able to predict which reactions should be carried out in a test tube in order to observe nonprecipitate results (e.g., heat from acid/base reactions, gas being given off, small amounts of precipitate
formed).
Copper Cycle:
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What was the purpose of adding nitric acid in the first step?
Why was an excess of NaOH necessary in the second step?
Why was the solution heated in the third step? Why was filtering necessary?
What role did H2SO4 play in the fourth step?
What did the zinc granules do the reaction in the fifth step? What happened if not enough zinc was
added?
Identify how experimental errors in each step would affect your overall results.
Be able to calculate the percent recovered at the end of the reaction.
Limiting Reactant:
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Be able to calculate the theoretical yield of a product given starting amounts of two reagents. Use
stoichiometry to determine which reagent will yield the least amount (moles or grams) of solid product.
Why were the test tubes heated in a hot water bath? Why did texture/appearance of the precipitates
vary?
What was the precipitate formed in this experiment?
Why did the height of precipitate level off from trials 4 to 6?
Describe how you would test the supernatant solution in a reaction. Be able to identify the ions present
in a supernatant solution.
Be able to draw a stick donkey.
Ascorbic Acid Titration:
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What was the purpose of rinsing and conditioning the buret and pipet?
Why should you not force the last drop out of the pipet?
Why was the Vitamin C tablet heated? Why was the solution cloudy after heating?
What does it mean to standardize a solution?
Why should the titrant be added slowly in the first trial of a titration?
Why did the solution flash pink and then clear early in the titration? Why does the solution remain
pink at the endpoint of a titration? Identify the indicator used in acid/base reactions. What species
must be present in solution to make phenolphthalein clear or pink?
Be able to identify the ions present in solution before and after the endpoint.
Be able to write the balanced chemical equation for the reaction between an acid and a base.
Calculate concentration of unknown solution
Calculate mass using solution stoichiometry
Be able to identify sources of error of experimental techniques. Identify how those errors will affect
the calculated concentration of an acid or a base.
Gas Laws:
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Be able to predict the products of gas-forming reactions and balance the equation.
Use stoichiometry to calculate moles of gas predicted to form from a given amount of starting material.
Use the Ideal Gas Law to calculate the predicted volume of gas that should be produced from the moles
calculated above.
Given vapor pressure of water at a specific temperature and atmospheric pressure, calculate the pressure
due to the gas produced in a reaction (Dalton’s law of partial pressures)
What was the purpose of the leveling bulb in the experiment? Why did it have to be lowered as the
reaction proceeded? What happens to the levels of water in the buret if the reaction flask is removed?
Given the volume of gas produced in a reaction, the pressure of gas, and room temperature, calculate
the number of moles of gas generated in a reaction.
Be able to describe experimental errors that result in a high or low percent yield.
Determine how experimental errors would affect the percent yield.
Thermochemistry:
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Describe the experimental set-up for the reactions carried out in this experiment. Why were Styrofoam
cups used in the reactions? What is the purpose of a closed system in this experiment?
Calculate the amount of heat given off or absorbed during a reaction using q = msΔT, where q is heat in
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Joules, m is mass in g, s is specific heat in J/g· C, and ΔT is Tf – Ti.
Why were heats of reactions calculated per mole of water formed?
Why were linear regressions only performed on the downward sloping portions of the graphs?
How was the final temperature determined from the graphs?
Use Hess’ Law to calculate the enthalpy of a reaction.