EXPANDED COURSE OBJECTIVES: CH115 General Chemistry I

Course Objectives: Upon completion of this course, the student should be able to:

Ø Offer examples of how chemistry is used in everyday life.

Ø Set forth what the scientific method is and how it is used.

Ø Classify materials in terms of homogeneous and heterogeneous mixtures.

Ø Distinguish between compounds and elements.

Ø Compare physical versus chemical properties.

Ø Recall from memory commonly used prefixes used with SI units.

Ø Solve problems involving density, volume, and mass.

Ø Convert between the Kelvin, Celsius, and Fahrenheit temperature scales.

Ø Apply scientific notation and correct number of significant figures in problem solving.

Ø Discuss the difference between accuracy and precision.

Ø Utilize factor-label method of problem solving.

Ø Recall from memory common conversion factors for the metric system to English system. Examples of such conversion factors include grams to pounds, centimeters to inches and liters to gallons.

Ø Restate the points of Dalton’s atomic theory.

Ø Distinguish between the law of definite proportions and the law of multiple proportions.

Ø Argue how Dalton’s third hypothesis is another way of stating the law of conservation of mass.

Ø Explain how electrons were discovered and how Millikan’s oil drop experiment determined the charge of the electron.

Ø Predict the path of alpha particles, beta particles, and gamma rays as they pass between two oppositely charged electrical plates.

Ø Set forth how Rutherford’s experiment concluded that atoms are mostly empty space with very small central cores, which are known as nuclei.

Ø Predict the path of protons, electrons, and neutrons as they pass between oppositely charged electrical plates.

Ø Compute the number of electrons, protons, and neutrons in atoms and ions.

Ø Give examples of isotopes.

Ø Recall from memory the names for the three isotopes of hydrogen.

Ø Predict if an element is a metal, nonmetal, or metalloid.

Ø Classify elements as alkali metals, alkaline earth metals, or noble gases.

Ø List several examples of diatomic molecules.

Ø Classify ions in terms of momatomic ions, polyatomic ions, cations, and anions.

Ø Distinguish between molecular and empirical formulas.

Ø Predict correct formulas for ionic compounds.

Ø Name common ionic compounds, molecular compounds, binary acids, oxoacids, bases, and hydrates given their respective chemical formulas.

Ø Predict the chemical formulas of common ionic compounds, molecular compounds, binary acids, oxoacids, bases, and hydrates given their respective names.

Ø Convert between grams and atomic mass units (AMU’s).

Ø Calculate average atomic mass given the mass and natural abundance of each isotope.

Ø Recall from memory Avogadro’s number.

Ø Determine the number of objects present in a given number of moles.

Ø Convert between mass, number of moles, and number of atoms (molecules) of an element (compound).

Ø Establish the molecular mass and molar mass given the molecular formula.

Ø Sketch the main components of a mass spectrometer and comment on its use.

Ø Compute the percent composition (mass percent) given the chemical formula for an ionic or molecular compound.

Ø Describe the experimental procedure used to determine empirical formulas.

Ø Establish the molecular formula given the mass of each element present (or mass percent of each element) and the compound’s molar mass.

Ø Balance chemical equations. (see CSB Homework in section 5.4 as well)

Ø Interpret the meaning of chemical equations in terms of molecules, moles, and masses.

Ø Distinguish between products and reactants in a chemical equation.

Ø Predict the products formed by combustion reactions.

Ø Use stoichiometric methods to predict the mass (number of moles) of the products formed given the mass of each reactant (number of moles of each reactant).

Ø Use stoichiometric methods to deduce the limiting reagent, excess reagent, the amount of expected products produced, and the amount of excess reagent left over upon completion of the reaction given the mass (number of moles) of each reactant in the chemical equation.

Ø Use stoichiometric methods to predict the theoretical yield and percent yield given the mass (number of moles) of each reactant and the actual yield of a reaction.

Ø Calculate the mass (number of moles) of each reactant required given the percent yield and the mass (number of moles) of products desired.

Ø Distinguish between solute, solvent, and solution.

Ø Classify common compounds as strong electrolytes, weak electrolytes, or nonelectrolytes (strong or weak acids or strong or weak bases).

Ø Suggest why water is often called a universal solvent utilizing the terms polar solvent, dissociation, ionization, and hydration.

Ø Describe precipitation reactions using the terms solubility and precipitate.

Ø Classify common ionic compounds as soluble or insoluble.

Ø Predict the resulting products and write the molecular equation, ionic equation, and net ionic equation and identify spectator ions given the reactants of a chemical reaction.

Ø Distinguish between Arrhenius acids and bases and Brfnsted acids and bases.

Ø Compare and contrast the properties of acids and bases.

Ø List common examples of monoprotic, diprotic and triprotic acids.

Ø Justify how some ions can act as an acid or as a base (amphoteric).

Ø Explain, by using a chemical equation, how ammonia (NH₃) is classified as a Brfnsted base.

Ø Predict the products formed by acid-base neutralization reactions.

Ø Discuss what factor results in an oxidation-reduction reaction.

Ø Identify oxidation half-reactions, reduction half-reactions, oxidizing agents and reducing agents.

Ø Assign oxidation numbers to elements in compounds and ions.

Ø Categorize redox reactions in terms of combination reactions, decomposition reactions, displacement reactions, and disproportionation reactions.

Ø Predict the results of a chemical reaction involving metals given the activity series (electrochemical series).

Ø Predict the results of a chemical reaction involving halogens given the halogen activity series.

Ø Compute the molarity of a solution given the mass (number of moles) of solute and the volume of solution.

Ø Describe the method for preparing a specific molar solution given the volume of solution required and the solute to be used.

Ø Relate in detail how to prepare a specific dilute solution given a known stock solution using dilution techniques.

Ø Predict the mass of a precipitate formed using gravimetric analysis methods.

Ø Deduce the mass percent of specific ions present in an original solution given the results of a gravimetric analysis.

Ø Use the terms titration, standard solution, equivalence point, and indicator to describe quantitative studies of acid-base neutralization reactions.

Ø Determine the concentration of an unknown acid (base) given the results of an acid-base titration.

Ø Predict the amount (mass, moles, or volume of solution) of an acid (base) required to neutralize a base (acid).

Ø Predict the volume of an oxidizing (reducing) agent solution required to oxidize (reduce) a specific volume of reducing (oxidizing) agent solution provided that the net ionic equation is given.

Ø Recall from memory at least ten common substances that are gases at one atmosphere and 25^o C.

Ø Distinguish between the terms gas and vapor.

Ø List four physical characteristics of all gases.

Ø Define the terms velocity, acceleration, force, newton, energy, joule, kinetic energy, pressure and pascal.

Ø Describe how a simple barometer is constructed and how it functions.

Ø Convert between torr, mmHg, atmospheres, and pascals.

Ø State the difference between open-tube manometers and closed-tube manometers and indicate how each is used.

Ø Write, explain, and apply each of the following:

Ø Boyle’s law (P µ 1/ V and P₁V₁ = P₂V₂).

Ø Charles’ law (P µ T and V₁/ T₁ = V₂/ T₂).

Ø Avogadro’s law (V µ n).

Ø Ideal gas law (PV = nRT).

Ø Describe the Kelvin temperature scale.

Ø Recall from memory the gas constant (R).

Ø State what standard temperature and pressure (STP) are and demonstrate that at STP one mole of gas occupies 22.4 liters.

Ø Perform calculations involving density, the Ideal gas equation and molar mass.

Ø Use the Ideal gas equation to determine the moles of a gas and use the number of moles in stoichiometric-based problems.

Ø State Dalton’s law of partial pressures and utilize it in problems involving mixtures of gases including the collection of gases over water.

Ø Define mole fraction and verify that P_a = X_aP_total.

Ø Discuss the four assumptions upon which the kinetic molecular theory of gases is based.

Ø Suggest how the kinetic molecular theory of gases qualitatively explains the following:
The compressibility of gases .Boyle’s law. Charles’ law. Avogadro’s law. Dalton’s law of partial pressure.

Ø Perform calculations using root-mean-square speeds.

Ø Describe the process of gaseous diffusion.

Ø Argue how a real gas, behaving non-ideally, differs from an ideal gas as described by the four assumptions in the kinetic molecular theory of gases.

Ø Conclude under what conditions a real gas will approximate an ideal gas.

Ø Apply van der Waal’s equation to real gases.

Ø Define and explain the following terms:
Energy, Radiant energy, Thermal energy, Chemical energy, Potential energy, Thermochemistry, Open system, Closed system, Isolated system, Endothermic, Exothermic, Enthalpy (DH), Calorimetry, Heat capacity, Specific heat,

Ø Classify common processes as endothermic or exothermic.

Ø Use thermochemical equations and stoichiometry to determine amount of heat lost or gained in a chemical reaction.

Ø Perform calculations involving specific heat, mass and temperature change.

Ø Sketch the main components of a constant-volume bomb calorimeter.

Ø Determine heats of reactions given experimental data collected in a calorimetry experiment.

Ø Calculate standard enthalpy of reactions given the standard enthalpy of formations for products and reactants.

Ø Apply Hess’s law to a multi-step process to determine standard enthalpy of reaction.

Ø Describe heat of solution, lattice energy, heat of hydration, heat of dilution, system, surrounding, and internal energy.

Ø Classify properties of materials as state functions or non-state functions.

Ø Restate the First Law of Thermodynamics.

Ø Recall the sign conventions for work and heat used in the textbook.

Ø Apply heat and work relationships to gas-phase problems.

Ø Define H in term of E, P, and V.

Ø Calculate change in internal energy (DE) given thermochemical equations.

Ø Recall Explain how Planck’s theory challenged classical physics.

Ø Define wavelength, frequency, and amplitude of waves.

Ø Utilize the relationship between speed, wavelength, and frequency (hertz).

Ø Describe Maxwell’s theory of electromagnetic radiation.

Ø Recall from memory the speed of light (3.00 x 10⁸ m/s).

Ø Apply the metric unit of nano in calculations involving wavelength of light.

Ø Classify various regions of the electromagnetic spectrum in terms of energy, frequency and wavelength.

Ø Use Planck’s equation to determine energy, frequency, or wavelength of electromagnetic radiation.

Ø Describe the photoelectric effect as explained by Einstein using such terms as threshold frequency, photons, kinetic energy, binding energy, light intensity and number of electrons emitted.

Ø Show how Bohr’s model of the atom explains emission, absorption and line spectra for the hydrogen atom.

Ø Compare Bohr’s model of the atom and that of the sun and surrounding planets.

Ø Predict the wavelength (frequency) of electromagnetic radiation emitted (absorbed) for electronic transitions in a hydrogen atom.

Ø Use the terms ground state and excited state to describe electronic transitions.

Ø Describe De Broglie’s relationship involving the wavelength of particles.

Ø Explain why for common objects traveling at reasonable speeds the corresponding wavelength becomes vanishingly small.

Ø Explain the major components of a laser and list three properties that are characteristic of a laser.

Ø Describe Heisenberg’s uncertainty principle.

Ø Contrast orbits (shells) in Bohr’s theory with orbitals in quantum theory.

Ø Discuss the concept of electron density.

Ø Recall from memory the four quantum numbers (n, ℓ, m_ℓ, m_s) and their relationships.

Ø Relate the values of the angular momentum quantum number, ℓ, to common names for each orbital (s, p, d, f) and describe their shapes.

Ø Account for the number of orbitals and number of electrons associated with each value of ℓ, the angular momentum quantum number.

Ø Categorize orbital energy levels in many-electron atoms in order of increasing energy.

Ø Write the four quantum numbers for all electrons in multi-electron atoms.

Ø Predict the electron configuration and orbital diagrams for multi-electron atoms using the Pauli exclusion principle and Hund’s rule.

Ø Deduce orbital diagrams from diamagnetic and paramagnetic data.

Ø Derive the ground state electron configuration of multi-electron atoms using the Aufbau principle.

Ø List several exceptions to the expected electron configuration for common metals (Cr, Mo, Cu, Ag, and Au).

Ø Explain the basis of the periodic table as described by Mendeleev and Meyer and indicate the shortcomings of their method.

Ø Explain the basis of the periodic table as described by Moseley and how it predicted properties of “missing” elements.

Ø Identify elements that correspond to each of the following groups: representative elements, noble gases, transition metals, lanthanides, actinides

Ø Describe the electron configuration of cations and anions and identity ions and atoms that are isoelectronic.

Ø Apply the concept of effective nuclear charge and shielding constants (screening constants) to justify why the first ionization energy is always smaller than the second ionization energy of a given atom.

Ø Predict the trends from left to right and top to bottom of the periodic table for each of the following: atomic radius, ionic radius, ionization energy, electron affinity, metallic character

Ø Relate why hydrogen could be placed in a class by itself when reviewing its chemical properties.

Ø Provide examples of Group 1A elements reacting with oxygen to form oxides, peroxides, and superoxides.

Ø Predict the reaction of alkali metals with water.

Ø Describe the reactivity of alkaline earth metals with water.

Ø Relate how strontium-90 could lead to human illness.

Ø Compare the reactivity of boron, a metalloid, to aluminum.

Ø Identify the metals, nonmetal, and metalloids of Group 4A.

Ø Recall the reactions that form nitric acid, phosphoric acid and sulfuric acid.

Ø List the halides (halogens)

Ø Indicate the three hydrohalic acids that are strong acids and the one hydrohalic acid that is a weak acid.

Ø Explain why the name for Group 8A has changed from inert gases to noble gases.

Ø List the three “coinage” metals and explain their relative inertness.

Ø Rationalize the characteristics of the properties of oxides of the third period elements.

Ø Classify oxides as acidic, basic, or amphoteric.

Ø Explain why concentrated bases such as NaOH should not be stored in glass containers

Ø Identify the valence electrons for all representative elements.

Ø Rationalize why alkali metals and alkaline earth metals usually form cations and oxygen and the halogens usually form anions using Lewis dot symbols in the discussion.

Ø Use Lewis dot symbols to show the formation of both ionic and molecular compounds.

Ø Define lattice energy, Coulomb’s law and the Born-Haber cycle.

Ø Demonstrate how the Born-Haber cycle is an application of Hess’s law and use the Born-Haber cycle to determine lattice energy for an ionic solid.

Ø Identify covalent compounds, the type of covalent bonds present, and the number of lone pairs of electrons using Lewis structures.

Ø Relate types of bonds to bond length and bond strength.

Ø Compare and contrast various properties expected for ionic compounds versus covalent compounds.

Ø Identify ionic, polar covalent and (nonpolar) covalent bonds using the concepts of electronegativity.

Ø Predict the relative changes in electronegativity with respect to position on the periodic table.

Ø Use the concept of electronegativity to rationalize oxidation numbers.

Ø Use Lewis dot and the octet rule to write Lewis structures of compounds and ions.

Ø Apply the concept of formal charge to predict the most likely Lewis structure of a compound.

Ø Explain how Lewis structures are inadequate to explain observed bond length (bond types) in some compounds and how the concept of resonance must be invoked.

Ø Recall several common examples in which the octet rule fails.

Ø Demonstrate, using Lewis structures, the formation of a coordinate covalent (dative) bond.

Ø Use Lewis structures and bond energies to predict heats of reaction.

Ø Rationalize why enthalpy change for breaking chemical bonds is positive and the formation of chemical bonds is negative.

Ø For each category identify the Molecular Geometry, Angle(s) sketch the shape

Category: AB₆E₀, AB₅E₁, AB₄E_2, AB₃E_3, AB₂E_4, AB₅E₀ AB₄E₁ AB₃E₂ AB₂E₃ AB₄E₀, AB₃E₁, AB₂E₂, AB₃E_0, AB₂E₁, AB₂E₀

Ø Identify using the VSEPR model, what category (and thus the corresponding molecular geometry, angle(s) and sketch) a molecular or ion belongs given its formula.

Ø Rationalize the observed decrease in angles for AB₄E₀, AB₃E₁, and AB₂E₂ and for AB₃E₀ and AB₂E₁.

Ø Apply VSEPR model to compounds with more than one central atom.

Ø Use the concepts of electronegativity, dipole moments, and VSEPR geometries to identify polar and nonpolar molecules.

Ø Use dipole moment concepts to predict properties of cis and trans isomers.

Ø Relate how a microwave oven functions and how the type of chemical bonds present effects the amount of energy absorbed.

Ø Sketch and justify how potential energy changes versus the interatomic distance for a diatomic molecule.

Ø Use Valence Bond theory, hybrid orbitals, and hybridization to explain the geometries predicted by VSEPR model.

Ø Identify what type of hybrid orbitals are in common compounds and ions.

Ø Apply the concepts of sigma and pi bonds and Valence Bond theory to explain properties of double and triple bonds and the concept of resonance.

Ø State what physical property of oxygen gas is not accounted for by Valence Bond theory but is in Molecular Orbital theory.

Ø Explain the difference between bonding and anti-bonding orbitals using the concepts of constructive and destructive interference of waves.

Ø Show molecular orbital energy diagrams for first and second row diatomic molecules identifying sigma and pi bonding and anti-bonding molecular orbitals.

Ø Write molecular orbital electron configurations for simple diatomics.

Ø Relate molecular orbital energy diagrams to bond order, bond length and bond strength.

Ø Describe resonance using the Molecular Orbital theory.

Ø Suggest the significance of the discoveries of fullerenes and nonotubes.