The Chemistry 1210 Final Exam consists of 40 questions and covers Chapters 1-10 and 12 from the 12th Edition of “Chemistry the Central Science” by Brown, LeMay, Bursten, Murphy, and Woodward.
To assist you on the exam you will be given the following information: CHEM 1210 AU13 Exam Supplemental Information Packet.
The following Practice Exams can serve as a starting point for your studying, but as you saw from the previous exams, the content will be the same, but the questions will be worded differently.
Dr. Zellmer is teaching CHEM 1250 this semester and his first three exams cover the chapters on the final. The links to his web-site are shown below:
The Following topics will be fair game on the exam:
Use significant figures, scientific notation, SI units and dimensional analysis in calculations.
Know common metric unit and SI units and prefixes, and be able to convert between units in calculations.
Empirical Information you are expected to know:
The volume relationship 1 cm3 = 1 mL
The meaning of the prefixes: kilo-, deci-, centi-, and milli- (Table 1.4)
Empirical Information you are not expected to memorize:
Conversions between metric and english units.
Prefixes other than those listed above.
Lab #1: Scientific Measurements
Know definition of density and be able to calculate quantities using it.
Know difference between accuracy and precision, and be able to distinguish between determinate and indeterminate errors and be able to apply these terms when analyzing a data set.
Lab #2: Separation of the Components of a Mixture
Distinguish among elements, compounds and mixtures (including separation methods for mixtures).
Distinguish between molecular substances and ionic substances in terms of their composition.
Describe the organization of the periodic table including the locations of metals and nonmetals.
Describe the structure of the atom in terms of protons, neutrons, and electrons.
Describe the electric charge and relative masses of protons, neutrons, and electrons.
Express the subatomic composition of isotopes/ions using chemical symbols together with atomic number and mass number.
Write the complete atomic symbol for ions (including atomic number, mass number, and charge).
Relate atomic weights to the masses of individual atoms and to their natural abundances.
Avogadro’s Constant and the Mole
Convert grams to moles and moles to grams using molar masses, and convert number of molecules to moles and moles to number of atoms or molecules using Avogadro’s number.
Common Cations and Anions
Know all the ions and charges on Table 2.4 and 2.5
Use the periodic table to predict the charges of common ions.
Write the empirical formulas of ionic compounds, given the charges of their component ions.
Write the name of an ion given its chemical formula, or write the chemical formula given its name.
Write the name of an ionic compound given its chemical formula, or write the chemical formula given its name.
Name or write chemical formulas for binary inorganic compounds and for acids.
Balance chemical equations.
Chemical and Physical Changes
Distinguish between chemical and physical change.
Identify simple combination, decomposition, and combustion reactions, predict their products, and write their balanced chemical equations.
Law of Conservation of Mass in chemical reactions.
Limiting Reactants and Theoretical Yield
Determine the limiting reactant in a reaction and use it to determine the amounts of products formed.
Calculate the percent yield of a reaction.
Empirical and Molecular Formulas
Calculate the empirical and molecular formula of a compound from percentage composition and molecular weight.
Quantitative Information From Chemical Equations
Be able to relate particle level diagrams to a balanced chemical equation and vice-versa.
Classify substances as either strong electrolytes, weak electrolytes or non-electrolytes. Demonstrate an understanding of the differences between the three.
Recognize and differentiate between strong acids, weak acids, strong bases and weak bases.
Reactions in Aqueous Solution
Identify simple acid-base, precipitation and redox reactions and be able to predict the products of such reactions.
Interpreting the Solubility Guidelines
You do not have to memorize these. They will be attached to the exam.
Net Ionic Equations
Given the molecular equation for a reaction be able to identify spectator ions and write the full ionic equation and/or the net ionic equation for the same reaction.
Assign oxidation numbers to individual atoms in neutral substances and ions, and use these assignments to determine which substance is reduced which substance oxidized in a redox reaction.
Interpreting the Activity Series
You do not have to memorize this series. It will be attached to the exam.
Reactions Involving Gaseous Products
Be able to predict which reactions will produce a gaseous product.
Calculate the molarity of a solution, and be able to convert between molarity, the number of moles present in a solution, and the volume of the solution.
Know how to prepare a dilute solution with a specific concentration and volume from a more concentrated solution.
Stoichiometry in Solution and Lab #3: Development of an Equation
Determine limiting reactants and/or calculate theoretical yields of reactions involving aqueous solutions. Use the results of a titration to determine the concentration of an unknown solution.
Know what an equivalence point/end point of a titration is and be able to use it to perform calculations based on titration data.
Heat and Change in Enthalpy
Express the relationships among the quantities heat and change in ethalpy. Learn their sign conventions, including how the signs of heat and change in enthalpy relate to whether a process is exothermic or endothermic.
The First Law of Thermodynamics
State the first law of thermodynamics and understand how it is applied in chemical reactions.
Understand the concept of a state function and be able to give examples of quantities that are and are not state functions.
Stoichiometry of Thermochemical Reactions
Use thermochemical equations to relate the change in enthalpy to the amount of substance involved in the reaction.
Relate temperature measurements and heat transferred by using heat capacities or specific heats.
Use Hess’s law to determine enthalpy changes for chemical reactions.
Experiment #6: Calorimetry and Hess’s Law
Standard Enthalpy of Formation
Use standard enthalpies of formation to calculate the standard enthalpy change for reactions.
Explain what photons are, and be able to calculate their energies given either their frequency or wavelength. Calculate the wavelength of electromagnetic radiation given its frequency or its frequency given its wavelength. Order the common kinds of radiation in the electromagnetic spectrum according to their wavelengths or energy.
Quantized Energy and Photons
Use quantum theory to understand that energy is quantized and to explain the photoelectric effect.
The Bohr Hydrogen Atom
Using the Bohr theory, explain how line spectra relate to the idea of quantized energy states of electrons in atoms. Be able to identify the limitations of the Bohr Model.
Experiment #7: Emission of Light and Atomic Models
Relate the quantum numbers to the number and type of orbitals, and recognize the different orbital shapes.
Interpret radial probability function graphs for the orbitals. Draw an energy-level diagram for the orbitals in a many-electron atom, and describe how electrons populate the orbitals in the ground-state of an atom, using the Pauli Exclusion Principle and Hund’s rule.
Electron Configuration/Orbital Block Notation
Use the periodic table to write condensed electron configurations and determine the number of unpaired electrons in an atom.
Effective Nuclear Charge
Understand the meaning of effective nuclear charge and how the effective nuclear charge depends upon nuclear charge and electron configuration.
Use the periodic table to predict the trends in atomic radii, ionic radii, ionization energy, and electron affinity.
Explain how the radius of an atom changes upon losing electrons to form a cation or gaining electrons to form an anion.
Explain how the ionization energy changes as we remove successive electrons. Recognize the jump in ionization energy that occurs when the ionization corresponds to removing a core electron.
Electron Configuration of Ions
Be able to write the electron configurations of ions.
Understand how irregularities in the periodic trends for electron affinity can be related to electron configuration.
Be able to understand how lattice energy is dependent upon the charge and size of ions.
Be able to recognize how lattice energy relates to physical properties such as melting point and boiling point.
Understand qualitatively how the radii of anions and cations differ from neutral atoms
Be able to draw Lewis dot structures
Bond Polarity and Electronegativity
Know the trends in electronegativity and use them to distinguish between polar covalent bonds and non-polar covalent bonds (This PhET tutorial shown in class should help with this concept: http://phet.colorado.edu/en/simulation/molecule-polarity)
Be able to assign formal charges to all of the atoms on a Lewis dot structure
Use formal charges to predict the most stable Lewis Structure
Understand the concept and meaning of resonance structures and identify molecules where the bonding is best represented by resonance structures
Exceptions to the Octet Rule
Know the exceptions to the octet rule and how to draw Lewis structures where the octet rule is violated
Strengths of Covalent Bonds
Know the relationship between bond enthalpy and bond length and be able to estimate the ΔH of a reaction given the bond enthalpies
Understand the relationship between bond order (single, double, triple bonds), bond length, and bond strength
THe VSEPR Model
Use the VSEPR model to predict molecular geometries and electron domain geometries of molecules (This PhET tutorial shown in class should help with this concept: http://phet.colorado.edu/en/simulation/molecule-shapes)
Be able to determine the number of bonding pairs and the number of lone pairs surrounding a central atom in a molecule
Understand the bond angle trends (deviations from the ideal bond angles) in molecules with unshared pairs or multiple bonds on the central atom
Molecular Shape and Polarity
Be able to identify polar and non-polar molecules
Be able to identify the hybridization on the central atom in a molecule
Know the similarities and differences between the sp, sp2, and sp3 hybrid orbitals particularly regarding how many π bonds can be formed on each central atom
Be familiar with the term resonance as it relates to the bonding in a molecule
Be able to determine the number of sigma and pi bonds in a molecule
Know the orbital character of single, double, and triple bonds
Be familiar with the terms resonance, delocalization, and π bonding
Be familiar with the relationship between potential energy and the distance between atoms
Know how orbital overlap between atomic orbitals relates to the energies of the molecular orbitals in a molecular orbital diagram
Be able to identify and distinguish between σ, σ*, π, and π* interactions in a molecule
Be able to use a molecular orbital diagram to determine the bond order in a molecule or ion
Be able to distinguish between the terms paramagnetic and diamagnetic
Be able to use a molecular orbital diagram to determine if a molecule is paramagnetic or diamagnetic
Recognize Pressure is force per unit area. Know the common units of pressure such as pascals, bar, atmosphere, and torr and be able to convert between units in calculations.
Understand how a barometer is used to measure atmospheric pressure and how a manometer is used to measure the pressure of an enclosed gas.
Understand the relationships given in Boyle’s, Charles, and Avogadro’s Law and how each of these laws is a special case of the ideal-gas equation.
Ideal Gas Law
Use the ideal gas equation to calculate variations in pressure, volume, number of moles, and temperature when one or more of the others is changed.
Know the conditions at STP.
Further Applications of the Ideal Gas Law
Use the ideal-gas equation to relate the density of a gas and its molar mass, and then to calculate the volumes of gases formed or consumed in a chemical reaction.
Gas Mixtures and Partial Pressures
Recognize the total pressure of a gas is the sum that each gas would exert if it were present alone (under the same conditions).
Be able to calculate the total pressure of a gas mixture given its partial pressures, or given information for calculating partial pressures.
Know the relationship between partial pressure and mole fraction.
Kinetic Molecular Theory of Gases
Describe the kinetic-molecular theory of gases and how it explains the pressure and temperature of a gas, the gas laws, and the rates of effusion and diffusion.
Understand that individual molecules of a gas do not all have the same kinetic energy at a given instant and that their speeds are distributed over a wide range, and be able to identify the characteristics of a gas that influence this distribution.
Molecular Effusion and Diffusion
Differentiate between effusion and diffusion and understand how the mean free path of gas molecules influence the rate of effusion and diffusion.
Explain why intermolecular attractions and molecular volumes cause real gases to deviate from ideal behavior at high pressure or low pressure.
Identify the intermolecular attractive interactions that exist between atoms, molecules or ions based on the composition and molecular structure of a substance, and compare the relative strengths of these interactions.
Select Properties of Liquids
Explain the concepts of viscosity and surface tension in liquids. Compare the relative strengths of viscosity and/or surface tension based on the composition and molecular structure.
Describe the differences in solids, liquids and gases and the characteristic properties of each. Know the names of the various changes in state between these three states of matter.
Interpret heating curves and be able to calculate quantities related to temperature and enthalpies of phase changes. Section 11.5
Explain how the vapor pressure of a liquid varies with changes in temperature and/or intermolecular forces. Describe the relationship between vapor pressure and boiling point and use it to predict the effect of changing pressure on the boiling point.
Be able to interpret and sketch phase diagrams. Define critical pressure, critical temperature, critical point, triple point, normal melting point and normal boiling point.
Classifications of Solids
Classify solids based on their bonding/intermolecular forces and describe the ways in which these forces relate to the physical properties of solids.
Structures of Solids
Explain the differences between crystalline and amorphous solids. Identify each of the four two-dimensional and seven three-dimensional primitive crystal lattices. Show the positions of the lattice points in body-centered and face-centered lattices, given a pattern of atoms in a crystalline solid identify the lattice vectors and unit cell.
Recognize common structures of metallic and ionic solids. Calculate the unit cell dimensions, empirical formula and density of these solids from a picture of the unit cell and tabulated values of ionic/atomic radii.
Be able to distinguish between hexagonal close packing and cubic close packing.
Describe the similarities and differences between the electron sea model and molecular orbital model of metallic bonding. Use molecular orbital theory to qualitatively predict periodic trends in melting point, boiling point, and hardness of metallic elements.
Predict the structures of ionic solids from their ionic radii and empirical formula and be able to determine the coordination number of the cation/anion in ionic solids.
Describe the basic electronic structure of a semiconductor, and use the periodic table to qualitatively compare the band gap energies of semiconductors.
Explain how doping can be used to control the conductivity of semiconductors and identify elements that can be used to n-dope or p-dope a semiconductor.