Download Chemistry H proficiencies

Chemistry H proficiencies
Topic
Chemistry: The
Study of Change
Atoms,
Molecules, Ions
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Mass
Relationships in
Chemical
Reactions
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Proficiencies
Set forth what the scientific method is and how it is used.
Classify materials in terms of homogeneous and heterogeneous mixtures.
Distinguish between compounds and elements.
Compare physical versus chemical properties.
Recall from memory commonly used prefixes used with SI units. (nano-, micro-, milli-, centi-, deci-,
kilo-)
Solve problems involving density, volume, and mass.
Convert between the Kelvin, Celsius, and Fahrenheit temperature scales.
Apply scientific notation and correct number of significant figures in problem solving.
Discuss the difference between accuracy and precision.
Utilize factor-label method of problem solving.
Recall the points of Dalton’s atomic theory.
Distinguish between the law of definite proportions and the law of multiple proportions.
Explain how electrons were discovered and how Millikan’s oil drop experiment determined the charge
of the electron.
Predict the path of alpha particles, beta particles, and gamma rays as they pass between two oppositely
charged electrical plates.
Set forth how Rutherford’s experiment concluded that atoms are mostly empty space with very small
central cores, which are known as nuclei.
Predict the path of protons, electrons, and neutrons as they pass between oppositely charged electrical
plates.
Compute the number of electrons, protons, and neutrons in atoms and ions.
Give examples of isotopes.
Predict if an element is a metal, nonmetal, or metalloid.
Classify elements as alkali metals, alkaline earth metals, or noble gases.
List several examples of diatomic molecules.
Classify ions in terms of momatomic ions, polyatomic ions, cations, and anions.
Distinguish between molecular and empirical formulas.
Predict correct formulas for ionic compounds.
Name common ionic compounds, molecular compounds, binary acids, oxoacids, bases, and hydrates
given their respective chemical formulas.
Predict the chemical formulas of common ionic compounds, molecular compounds,
binary acids, oxoacids, bases, and hydrates given their respective names.
Convert between grams and atomic mass units (AMU’s).
Calculate average atomic mass given the mass and natural abundance of each isotope.
Recall from memory Avogadro’s number.
Determine the number of objects present in a given number of moles.
Convert between mass, number of moles, and number of atoms (molecules) of an element
(compound).
Establish the molecular mass and molar mass given the molecular formula.
Compute the percent composition (mass percent) given the chemical formula for an ionic or molecular
compound.
Establish the molecular formula given the mass of each element present (or mass percent of each
element) and the compound’s molar mass.
Balance chemical equations.
Interpret the meaning of chemical equations in terms of molecules, moles, and masses.
Distinguish between products and reactants in a chemical equation.
Predict the products formed by combustion reactions.
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).
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
Reactions in
Aqueous
Solution
Gases
the mass (number of moles) of each reactant in the chemical equation.
15. 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.
16. Calculate the mass (number of moles) of each reactant required given the percent
yield and the mass (number of moles) of products desired.
1. Distinguish between solute, solvent, and solution.
2. Classify common compounds as strong electrolytes, weak electrolytes, or nonelectrolytes (strong or
weak acids or strong or weak bases).
3. Suggest why water is often called a universal solvent utilizing the terms polar solvent, dissociation,
ionization, and hydration.
4. Describe precipitation reactions using the terms solubility and precipitate.
5. Classify common ionic compounds as soluble or insoluble.
6. 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.
7. Distinguish between Arrhenius acids and bases and Bronsted acids and bases.
8. Compare and contrast the properties of acids and bases.
9. List common examples of monoprotic, diprotic and triprotic acids.
10. Justify how some ions can act as an acid or as a base (amphoteric).
11. Explain, by using a chemical equation, how ammonia (NH3) is classified as a Bronsted base.
12. Predict the products formed by acid-base neutralization reactions.
13. Discuss what factor results in an oxidation-reduction reaction.
14. Identify oxidation half-reactions, reduction half-reactions, oxidizing agents and reducing agents.
15. Assign oxidation numbers to elements in compounds and ions.
16. Categorize redox reactions in terms of combination reactions, decomposition reactions, displacement
reactions, and disproportionation reactions.
17. Predict the results of a chemical reaction involving metals given the activity series (electrochemical
series).
18. Predict the results of a chemical reaction involving halogens given the halogen activity series.
19. Compute the molarity of a solution given the mass (number of moles) of solute and the volume of
solution.
20. Describe the method for preparing a specific molar solution given the volume of solution required and
the solute to be used.
21. Relate in detail how to prepare a specific dilute solution given a known stock solution using dilution
techniques.
22. Predict the mass of a precipitate formed using gravimetric analysis methods.
23. Deduce the mass percent of specific ions present in an original solution given the results of a
gravimetric analysis.
24. Use the terms titration, standard solution, equivalence point, and indicator to describe quantitative
studies of acid-base neutralization reactions.
25. Determine the concentration of an unknown acid (base) given the results of an acid-base titration.
26. Predict the amount (mass, moles, or volume of solution) of an acid (base) required to neutralize a base
(acid).
27. 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.
1. List four physical characteristics of all gases.
2. Define the terms force, energy, joule, kinetic energy, pressure and pascal.
3. Describe how a simple barometer is constructed and how it functions.
4. Convert between torr, mmHg, atmospheres, and pascals.
5. State the difference between open-tube manometers and closed tube manometers and indicate how
each is used.
6. Write, explain, and apply each of the following:
Boyle’s law (P 1/ V and P1V1 = P2V2).
Charles’ law (P T and V1/ T1 = V2/ T2).
Avogadro’s law (V n).
Ideal gas law (PV = nRT).
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Thermochemistry
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Quantum Theory
and the
Electronic
Structure of
Atoms
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Describe the Kelvin temperature scale.
Recall from memory the gas constant R.
State what standard temperature and pressure (STP) are and demonstrate that at STP one mole of gas
occupies 22.4 liters.
Perform calculations involving density, the Ideal gas equation and molar mass.
Use the Ideal gas equation to determine the moles of a gas and use the number of moles in
stoichiometric-based problems.
State Dalton’s law of partial pressures and utilize it in problems involving mixtures of gases including
the collection of gases over water.
Define mole fraction and verify that Pa = xaPtotal.
Discuss the four assumptions upon which the kinetic molecular theory of gases is based.
Perform calculations using Graham’s Law.
Describe the process of gaseous diffusion.
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.
Conclude under what conditions a real gas will approximate an ideal gas.
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 (H),
Calorimetry, Heat capacity, Specific heat
Classify common processes as endothermic or exothermic.
Use thermochemical equations and stoichiometry to determine amount of heat lost or gained in a
chemical reaction.
Perform calculations involving specific heat, mass and temperature change.
Determine heats of reactions given experimental data collected in a calorimetry experiment.
Calculate standard enthalpy of reactions given the standard enthalpy of formations for products and
reactants.
Apply Hess’s law to a multi-step process to determine standard enthalpy of reaction.
Explain how Planck’s theory challenged classical physics.
Define wavelength, frequency, and amplitude of waves.
Utilize the relationship between speed, wavelength, and frequency (hertz).
Recall from memory the speed of light (3.00 x 10 8 m/s).
Apply the metric unit of nano in calculations involving wavelength of light.
Classify various regions of the electromagnetic spectrum in terms of energy, frequency and
wavelength.
Use Planck’s equation to determine energy, frequency, or wavelength of electromagnetic radiation.
Describe the photoelectric effect.
Show how Bohr’s model of the atom explains emission, absorption and line spectra for the hydrogen
atom. Use of Bohr’s equation is optional but not necessary.
Compare Bohr’s model of the atom and that of the sun and surrounding planets.
Use the terms ground state and excited state to describe electronic transitions.
Explain how matter can have wavelike properties i.e. De Broglie wavelength, and that this is
insignificant for normal sized objects.
Qualitatively describe Heisenberg’s uncertainty principle.
Contrast orbits (shells) in Bohr’s theory with orbitals in quantum theory.
Discuss the concept of electron density.
Recall from memory the four quantum numbers (n, l, ml, ms) and their relationships.
Relate the values of the angular momentum quantum number, l, to common names for each orbital (s,
p, d, f) and describe their shapes.
Account for the number of orbitals and number of electrons associated with each value of , the angular
momentum quantum number.
Categorize orbital energy levels in many-electron atoms in order of increasing energy.
Write the four quantum numbers for all electrons in multielectron atoms.
Predict the electron configuration and orbital diagrams for multi-electron atoms using the Pauli
exclusion principle and Hund’s rule.
Periodic
Relationships
Among the
Elements
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Chemical
Bonding – Basic
Concepts
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Chemical
Bonding –
Molecular
Geometry and
Hybridization of
Atomic Orbitals
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Deduce orbital diagrams from diamagnetic and paramagnetic data.
Derive the ground state electron configuration of multielectron atoms using the Aufbau principle.
List Cr and Cu as exceptions to the expected electron configuration for common metals.
Explain the basis of the periodic table as described by Mendeleev and indicate the shortcomings of his
method.
Explain the basis of the periodic table as described by Moseley and how it predicted properties of
“missing” elements.
Identify elements that correspond to each of the following groups:
representative elements,noble gases,transition metals, lanthanides,actinides
Describe the electron configuration of cations and anions and identity ions and atoms that are
isoelectronic.
Apply the concept of shielding to justify why the first ionization energy is always smaller than the
second ionization energy of a given atom.
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 (should be defined as < 0 for an
exothermic process),metallic character
Relate why hydrogen could be placed in a class by itself when reviewing its chemical properties.
Provide examples of Group 1A elements reacting with oxygen to form oxides, peroxides.
Predict the reaction of alkali metals with water.
Describe the reactivity of alkaline earth metals with water.
Compare the reactivity of boron, a metalloid, to aluminum.
Identify the metals, nonmetal, and metalloids of Group 14.
Recall the reactions that form nitric acid, phosphoric acid and sulfuric acid.
List the halides (halogens)
Explain HF is a weak acid while solutions of the other hydrogen halides are strongly acidic.
Rationalize the characteristics of the properties of oxides of the third period elements.
Classify oxides as acidic, basic, or amphoteric.
Identify the number of valence electrons for all representative elements.
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.
Use Lewis dot symbols to show the formation of both ionic and molecular compounds.
Identify covalent compounds, the type of covalent bonds present, and the number of lone pairs of
electrons using Lewis structures.
Relate types of bonds to bond length and bond strength.
Compare and contrast various properties expected for ionic compounds versus covalent compounds.
Identify ionic, polar covalent and (nonpolar) covalent bonds using the concepts of electronegativity.
Predict the relative changes in electronegativity with respect to position on the periodic table.
Use the concept of electronegativity to rationalize oxidation numbers.
Use Lewis dot and the octet rule to write Lewis structures of compounds and ions.
Apply the concept of formal charge to predict the most likely Lewis structure of a compound.
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.
Recall BeH2, BF3, NO, SF6, SCl2.
Use Lewis structures and bond energies to predict heats of reaction.
Rationalize why enthalpy change for breaking chemical bonds is positive and the formation of
chemical bonds is negative.
Fill in the following chart using the VSEPR model:
Category Molecular, Geometry, Angle(s) Sketch of the Shape
AB6E0, AB5E1, AB4E2, AB3E3, AB2E4, AB5E0, AB4E1, AB3E2, AB2E3, AB4E0, AB3E1, AB2E2,
AB3E0, AB2E1, AB2E0
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.
Rationalize the observed decrease in angles for AB 4E0, AB3E1, and AB2E2 and for AB3E0 and AB2E1.
Apply VSEPR model to compounds with more than one central atom.
Use the concepts of electronegativity, dipole moments, and VSEPR geometries to identify polar and
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Intermolecular
Forces and
Liquids and
Solids
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Physical
Properties of
Solution
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Chemical
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nonpolar molecules.
Sketch and justify how potential energy changes versus the interatomic distance for a diatomic
molecule.
Use Valence Bond theory, hybrid orbitals, and hybridization to explain the geometries predicted by
VSEPR model.
Identify what type of hybrid orbitals are in common compounds and ions.
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.
Characterize the properties of gases, liquids and solids in terms of density, compressibility and motion
of molecules.
Distinguish between intermolecular and intramolecular forces.
Identify and give examples of the following forces:
ion – ion, ion – dipole, ion – induced dipole, dipole – dipole,
dipole – induced dipole, induced dipole – induced dipole, van der Waals,
dispersion, hydrogen bonding
Suggest why H2O, HF and NH3 do not follow the expected trend as shown by the plot in Figure 11.7
Use the concepts of intermolecular forces to explain surface tension, capillary action, cohesion,
adhesion and viscosity.
Describe the structure of water and relate it to why ice is less dense then water.
Sketch density vs. temperature curve for water and relate how this plot has major significance for
aquatic life.
Distinguish between crystalline and amorphous solids and give examples of each.
Identify coordination numbers for atoms in simple cubic, face centered cubic and body-centered cubic
structures.
Determine the number of atoms contained in a unit cell for simple cubic, face centered cubic and
body-centered cubic structures.
Characterize ionic, covalent, molecular and metallic crystals including general properties and
examples of each.
Discuss the following:
evaporation, condensation, molar heat of vaporization, molar hear of fusion, boiling point,
melting point, supercooling, sublimation, deposition, molar heat of sublimation, phase changes,
critical temperature and pressure
Sketch a typical heating curve and identify various aspects of it.
Use phase diagrams to identify what phase(s) is/are present given specific conditions.
Use the terms saturated, unsaturated and supersaturated to describe solutions.
Distinguish between crystallization and precipitation.
Define, determine and inter-convert between each of the following:
molarity, percent by mass, mole fraction, molality
Suggest a shortcoming of molarity and explain why molality is a preferred concentration unit under
certain conditions.
Use the concept of fractional crystallization to show how dissolved solids can be separated.
State Henry’s law and use it to determine the solubility of gases in liquids.
Rationalize why two common materials (NH3 or CO2) when dissolved in water do not follow Henry’s
law.
Define colligative properties and give four examples (vapor –pressure lowering, freezing – point
lowering, boiling – point elevation).
Describe the apparatus used in fractional distillation.
Perform calculations involving boiling-point elevation, freezing-point depression, Kf, Kb and
molality.
Use the concepts of colligative properties to determine molar mass.
Define the van’t Hoff factor and demonstrate how it is incorporated into the colligative property
equations.
Give examples of common types of colloids and describe the dispersing medium and dispersed phase
for each.
Describe the Tyndall effect.
Distinguish between average and instantaneous rates of chemical reactions.
Kinetics
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Chemical
Equilibrium
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Acids and Bases
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Acids-Base
Equilibria and
Solubility
Equilibria
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Define rate constant.
Use the concepts of stoichiometry to write reaction rate expressions in terms of the disappearance of
reactants and the appearance of products.
Sketch the rate of reaction versus concentration of reactant for zero and first order reactions.
Use rate data to detemine rate laws and rate constants.
Show that half-life is independent of initial concentration of the reactant in a first order reaction.
Use the concept of half-life to determine concentration of reactants over time.
Describe the Collision Theory of Chemical Kinetics using the terms activation energy, activated
complex (transition state), potential energy profiles, endothermic and exothermic reactions.
Show that the sum of elementary steps is the overall reaction for a reaction mechanism and that
intermediates appear in the reaction mechanism but not in the overall reaction.
Define molecularity of unimolecular, bimolecular, and termolecular reactions.
Relate the importance of the rate-determining step in determination of reaction mechanisms.
Describe what a catalysis does, how it effects activation.
Relate the importance of enzymes as biological catalysts.
Describe chemical equilibrium using the terms forward and reverse reactions and dynamic process.
Write the equilibrium constant in terms of the equilibrium concentration of products and reactants and
their respective stoichiometric coefficients for both homogenous and heterogeneous equilibria.
Derive the relationship between Kp and Kc.
Determine equilibrium constant given equilibrium concentration data.
Relate equilibrium constant to rate constants from chemical kinetics.
Describe the relationship between reaction quotient and equilibrium constant and predict the direction
a reaction will proceed to reach equilibrium.
Use the concepts of equilibrium to determine concentration of all species in a solution.
Use Le Chatelier’s Principle to describe how changing concentration, volume, pressure, or
temperature will shift the reaction so that a equilibrium will be maintained.
Describe the effect of a catalyst has on equilibrium concentrations.
Compare and contrast Arrhenius, Brønsted, and Lewis acids and bases.
Describe what is meant by conjugate acid-base pairs and give several examples.
Use Kw to determine [H+] and [OH-] of solutions.
Discuss the pH scale and calculate pH and pOH given either [H+] or [OH-].
Define strong and weak acids and bases and give several examples of each.
Relate properties of conjugate acid-base pairs.
Determine Ka from experimental data.
Calculate pH, [H+] , weak acid concentration, and conjugate base concentration given Ka and the
initial concentration of the weak acid using the quadratic equation or the method of successive
approximation as needed.
Calculate percent ionization for a weak acid.
Calculate pH, [OH-], weak base concentration, and conjugate acid concentration given Kb and the
initial concentration of the weak base using the quadratic equation or the method of successive
approximation as needed.
Show the relationship between Ka,, Kb, and Kw.
Relate molecular structure and the strength of acids.
Predict the relative strengths of oxoacids.
Describe salt hydrolysis and explain how some salts produce neutral solutions, some acidic solutions
and others basic solutions.
Calculate the pH of salt solutions and determine the percent hydrolysis.
Describe that metal oxides form basic solutions with water, while nonmetal oxides form acidic
solutions with water.
Give several examples of Lewis acid-base reactions.
Describe the common ion effect as a special case of Le Châtelier’s principle.
Describe what a buffer solution is and its importance in chemical and biological systems.
Predict the pH profile of a strong acid-strong base titration and calculate the pH at any stage of the
titration.
Predict the pH profile of a strong acid-weak base (or strong base-weak acid) titration and calculate the
pH at any stage.
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6.
Entropy, Free
Energy, and
Equilibrium
Electrochemistry
Nuclear
Chemistry
Organic
Chemistry
Distinguish between end point and equivalence point of a titration.
Describe common acid-base indicators and suggest the correct method of selection for a specific
titration.
7. Use the concepts of equilibrium to relate ion product, Q, with Ksp to predict if a solution is
unsaturated, saturated or supersaturated.
8. Distinguish between solubility product, molar solubility and solubility.
9. Describe how changing pH can affect solubility.
10. Describe how solubility product principle is used in qualitative analysis.
1. State the first and second law of thermodynamics.
2. Provide several examples of spontaneous processes.
3. Give several examples of endothermic spontaneous processes.
4. Define entropy using the terms disorder or randomness.
5. Justify why entropy of one mole of steam is greater then the entropy of one mole of water.
6. Predict the sign on the change in entropy (S) for common processes.
7. Use thermodynamic tables to determine S reaction.
8. Express Gibbs free energy in terms of H, T and S.
9. Predict if a reaction is spontaneous, spontaneous in the reverse direction or at equilibrium from the
sign on G.
10. Use thermodynamic tables to determine G of a reaction.
11. Rationalize the direction of a spontaneous reaction given H, S and T.
1. Describe the concept of redox reactions using such terms as reduction, oxidation, reducing agents,
oxidizing agents and oxidation numbers.
2. Balance redox equations in both acidic and basic solutions.
3. Describe an electrochemical cell using such terms as oxidation, reduction, galvanic cell, electrolytic
cell, anode, cathode, half-cell reactions, salt bridge, cell voltage and emf.
4. Use standard cell diagrams to describe an electrochemical cell.
5. Use standard reduction potentials to predict the emf of a cell.
6. Use standard reduction potentials to justify the activity series for metals and the activity series for
halogens as described in section 4.4 of this textbook.
7. Define what Standard Hydrogen Electrode (SHE) means and relate its significance to the standard
reduction potential table.
8. Predict the outcome of reactions based on standard reduction potentials and Eooverall.
9. Predict the expected results from the electrolysis of molten sodium chloride, water and aqueous
sodium chloride.
1. Compare chemical reactions to nuclear reactions.
2. State the charge, mass and symbol for each of the following:
proton, neutron, electron or beta particle, positron, alpha particle
3. Balance nuclear equations using the concepts of conservation of mass and conservation of atomic
numbers.
4. Explain how nuclei having a neutron-to-proton ratio that is too high results in beta-particle emission
and nuclei having a neutron-to-proton ratio that is too low results in either positron emission or
electron capture.
5. Use the concepts of nuclear binding energy, mass defect and Einstein’s mass-energy equivalence
relationship to calculate nuclear binding energy per nucleon.
6. Describe a radioactive decay series using the terms parent and daughter.
7. Balance nuclear equations involving transmutation.
8. Compare nuclear fission to nuclear fusion and use the term thermonuclear reaction in the discussion.
1. Write the general formula for and provide examples for each of the following organic functional
groups:
a. alkane
b. alkene
c. alkyne
d. cycloalkane
e. aromatic
f. alcohol
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4.
g. ether
h. aldehyde
i. ketone
j. carboxylic acid
k. ester
l. amine
Use IUPAC rules to name
a. alkanes
b. alkenes
c. alkynes
Draw and name all of the possible isomers given the chemical formula of an C 5H12.
Use the following groups in naming organic compounds
a. methyl
b. ethyl
c. n-propyl
d. n-butyl
e. isopropyl
f. t-butyl
g. amino
h. fluoro
i. chloro
j. bromo
k. iodo
l. nitro
m. vinyl