Assessment Package

1. Assessment Philosophy

Assess what students can explain, represent, calculate, and transfer. A balanced assessment samples:

• Conceptual knowledge: definitions, relationships, and model limitations.
• Representational competence: particle diagrams, formulas, equations, graphs, and words.
• Quantitative reasoning: setup, units, magnitude, calculation, and significant figures.
• Experimental reasoning: variables, uncertainty, safety, evidence, and conclusions.
• Transfer: use of a familiar principle in a new surface context.

Avoid making an exam primarily a speed test. Permit adequate time, use readable notation, and reduce construct-irrelevant complexity. Publish a blueprint before each major assessment.

2. Mathematics and Chemistry Diagnostic

Administer ungraded during Week 1. Let students use the result to select support resources.

Questions

1. Write 0.00007250 in scientific notation.
2. Calculate (3.20 × 10^4)(2.5 × 10^-3) with appropriate significant figures.
3. Solve d = m/V for V.
4. Convert 45.0 cm to meters.
5. Convert 2.50 h to seconds.
6. If y = 4x, what happens to y when x triples?
7. If P = k/V, what plot should be linear: P versus V, or P versus 1/V?
8. A line passes through (2, 5) and (6, 17). Find its slope.
9. State the number of significant figures in 0.004050.
10. Add 12.11 + 0.3 + 4.005 using measurement-reporting rules.
11. What information is missing from the statement “the mass is 15”?
12. Classify water as element, compound, or mixture.
13. What subatomic particle determines the identity of an element?
14. What is the net charge of a particle with 11 protons and 10 electrons?
15. Balance __ H2 + __ O2 → __ H2O.
16. Which is larger: one mole of particles or one dozen particles?
17. Estimate without a calculator: is 498/2.01 closer to 25, 250, or 2500?
18. On a graph, what quantity belongs on the horizontal axis when time is deliberately varied?
19. Explain in one sentence why units should cancel in a conversion.
20. Name one study action that requires recall without looking at the answer.

Answer key and routing

1. 7.250 × 10^-5.
2. 8.0 × 10^1 or 80. with two significant figures.
3. V = m/d.
4. 0.450 m.
5. 9.00 × 10^3 s.
6. It triples.
7. P versus 1/V.
8. 3 in y-units per x-unit.
9. Four.
10. 16.4.
11. A unit and enough context to identify the quantity/sample.
12. Compound.
13. Proton count/atomic number.
14. +1.
15. 2, 1, 2.
16. One mole.
17. 250.
18. Horizontal/x-axis.
19. Valid conversion factors equal one and remove the starting unit while introducing the wanted unit.
20. Examples: blank-page recall, flashcards, practice question, or explaining from memory.

Suggested routing: 17–20 correct, begin ordinary coursework; 13–16, complete targeted Module 0 lessons; 0–12, complete Module 0 and meet math/chemistry support during Week 1. Do not use the diagnostic to exclude students.

3. Readiness Quiz Bank

These items can be randomized. Use three to five per module, allow retry, and provide explanatory feedback.

Module 1

1. Is density intensive or extensive? Intensive.
2. Which is a chemical property: color, melting point, density, or flammability? Flammability.
3. What is the SI base unit of mass? Kilogram.
4. How many significant figures are in 100.20? Five.
5. A sample has m = 18.2 g and V = 7.0 mL. Find density. 2.6 g/mL.

Module 2

1. What determines an element's identity? Number of protons.
2. How many neutrons are in ^23Na? 12.
3. Write the common ion formed by calcium. Ca2+.
4. Name FeO. Iron(II) oxide.
5. Write aluminum nitrate. Al(NO3)3.

Module 3

1. State the exact number of entities in a mole. 6.02214076 × 10^23.
2. What bridge connects grams to particle count? Moles.
3. What is the approximate molar mass of O2? 32.00 g/mol.
4. How many moles of O atoms are in one mole of CO2? 2 mol.
5. What does an empirical formula represent? Lowest whole-number atom ratio.

Module 4

1. May subscripts be changed to balance an equation? No.
2. What do coefficients represent? Relative numbers of entities and relative mole amounts.
3. What is the oxidation number of a free element? 0.
4. Is precipitation a phase-change or chemical-reaction category? Chemical reaction.
5. Balance N2 + H2 → NH3. N2 + 3 H2 → 2 NH3.

Module 5

1. What unit must be used with a coefficient ratio? Moles or entities.
2. What determines theoretical yield? The limiting reactant under the assumed reaction.
3. Define actual yield. Measured amount obtained.
4. Write the percent-yield equation. Actual/theoretical × 100%.
5. Can initial grams alone identify the limiting reactant? No.

Module 6

1. Define molarity. Moles solute per liter solution.
2. Does dilution change solute moles in the ideal transfer model? No.
3. Which conducts: ions fixed in a solid lattice or mobile ions in solution? Mobile ions in solution.
4. Should AgCl(s) be split in a complete ionic equation? No.
5. What is a spectator ion? An ion present but unchanged on both sides of the ionic equation.

Module 7

1. At equivalence, what relationship has been met? Balanced-reaction stoichiometry.
2. What species is reduced: oxidizing agent or reducing agent? Oxidizing agent.
3. Oxidation is an increase or decrease in oxidation number? Increase.
4. Write the strong acid–strong base net ionic equation. H+ + OH− → H2O.
5. Does M1V1 = M2V2 automatically cover a 1:2 reaction? No.

Module 8

1. Define heat. Energy transferred because of a temperature difference.
2. If heat enters the system, what is the sign of q? Positive.
3. Write ΔT. Tfinal − Tinitial.
4. What quantity does specific heat express? Energy per mass per temperature change.
5. In an insulated setup, what is the sum of energy changes? Approximately zero.

Module 9

1. What is the sign of ΔH for an exothermic process? Negative.
2. What happens to ΔH when coefficients double? It doubles.
3. State Hess's-law basis. Enthalpy is a state function.
4. What is ΔH°f of an element in its standard state? Zero.
5. Does a catalyst change ΔH? No.

Module 10

1. Relate wavelength and frequency. Inversely through c = λν.
2. Which has more photon energy: shorter or longer wavelength? Shorter.
3. How many orbitals are in a p subshell? Three.
4. What is the maximum electron capacity of one orbital? Two.
5. Is an orbital a fixed electron path? No.

Module 11

1. State Hund's rule in brief. Singly occupy degenerate orbitals before pairing.
2. Write Na in noble-gas notation. [Ne]3s1.
3. Are cations generally smaller or larger than their neutral atoms? Smaller.
4. What generally happens to atomic radius across a period? It decreases.
5. What generally happens to first ionization energy across a period? It increases, with exceptions.

Module 12

1. How many total valence electrons are in CO2? 16.
2. Write the formal-charge equation. Valence − nonbonding − half bonding.
3. Can resonance contributors differ in atom positions? No.
4. What is the net formal charge sum for NO3−? −1.
5. Which attraction generally grows with higher ionic charge? Lattice attraction/lattice-energy magnitude.

Module 13

1. How many domains does a double bond count as? One.
2. What is the molecular geometry of NH3? Trigonal pyramidal.
3. What is the electron-domain geometry of H2O? Tetrahedral.
4. Why is CO2 nonpolar overall? Its equal bond dipoles cancel in a linear geometry.
5. What bonds make up a double bond? One sigma and one pi.

Module 14

1. Why must gas-law temperature be in kelvins? Gas proportionalities use absolute temperature.
2. State Dalton's law. Total pressure is the sum of partial pressures.
3. Do equal-temperature gases have equal average translational kinetic energy? Yes.
4. Are dispersion forces present in all atoms and molecules? Yes.
5. During an ideal phase-change plateau, does temperature rise? No; the energy changes phase proportion/potential-energy relationships.

4. Unit Exam Blueprints

Unit Exam 1 — Quantifying Chemical Composition and Change

Suggested length: 75 minutes, 100 points.

Unit Exam 2 — Solutions and Energy

Unit Exam 3 — Atomic Structure and Bonding

Cumulative Final

Suggested length: 120 minutes, 150 points.

At least 35% of the final should require integration of two or more modules. Example: determine gas moles stoichiometrically, then apply PV=nRT; or infer geometry and polarity, then compare boiling behavior.

5. Cumulative Practice Bank

Conceptual and representational items

1. Distinguish a homogeneous mixture from a compound using composition and separation.
2. Explain why density can identify a material only under specified conditions and with uncertainty.
3. Draw particle models for an element, a molecular compound, an ionic solid, and a homogeneous mixture.
4. Explain why ^35Cl and ^37Cl have similar chemistry but different masses.
5. Describe what 3 CaCl2 communicates at particle and mole scales.
6. Explain why balanced equations conserve atoms but may not conserve gas volume.
7. Distinguish theoretical yield from an expected experimental result.
8. Draw a particle view before and after dissolving a strong electrolyte.
9. Explain why a weak electrolyte can dissolve yet conduct weakly.
10. Distinguish an endpoint from an equivalence point.
11. Explain the signs of q_solution and q_reaction when solution warms.
12. Explain why breaking bonds requires energy while some bond-forming reactions are exothermic overall.
13. Distinguish heat from temperature.
14. Explain why line spectra reject a model with continuously variable atomic energies.
15. Distinguish an orbital from an orbit.
16. Explain the atomic-radius trend across Period 3 using attraction and shielding.
17. Compare Na and Na+ radius using electron structure.
18. Explain resonance without suggesting that molecules switch between drawings.
19. Distinguish electron-domain geometry from molecular geometry.
20. Explain how a molecule with polar bonds can be nonpolar.
21. Draw particle-level representations of gas compression at constant temperature.
22. Explain why kelvins are required in gas proportionalities.
23. Compare average kinetic energy and average speed for He and Xe at the same temperature.
24. Explain why branching can lower boiling point among similar nonpolar molecules.
25. Explain the flat region of a heating curve without saying added heat disappears.

Quantitative and structured-response items

1. Convert 3.75 × 10^2 μm to meters.
2. A 15.625 g solid raises water volume from 12.2 to 14.0 mL. Find density with justified precision.
3. An isotope occurs at 72.0% with mass 62.930 u and at 28.0% with mass 64.928 u. Find average mass.
4. Write formulas for calcium phosphate, ammonium sulfite, copper(II) nitrate, and dinitrogen trioxide.
5. Name K2Cr2O7, Fe2O3, P4O10, and NaHCO3.
6. Convert 4.50 × 10^24 molecules of water to moles.
7. Find the number of oxygen atoms in 0.250 mol Al2(SO4)3.
8. Calculate percent oxygen by mass in CaCO3.
9. Determine the empirical formula for 52.14% C, 13.13% H, and 34.73% O.
10. Balance C2H6 + O2 → CO2 + H2O.
11. For 2 KClO3 → 2 KCl + 3 O2, find moles O2 from 0.800 mol KClO3.
12. For N2 + 3 H2 → 2 NH3, find limiting reactant and ammonia moles from 1.50 mol N2 and 3.60 mol H2.
13. If 8.50 g product is obtained from a theoretical 10.2 g, find percent yield.
14. Find molarity of 5.85 g NaCl in 500.0 mL solution.
15. Find volume of 6.00 M stock required for 250.0 mL of 0.400 M solution.
16. Write the net ionic equation for aqueous calcium chloride plus sodium carbonate.
17. Find moles Na+ in 35.0 mL of 0.250 M Na2SO4.
18. 30.00 mL HCl requires 24.60 mL of 0.1150 M NaOH. Find HCl molarity.
19. Find heat required to warm 80.0 g water by 22.5 °C.
20. A reaction releases 1.85 kJ into 100.0 g solution with c = 4.184 J g−1 °C−1. Estimate ΔT if calorimeter heat is neglected.
21. If a thermochemical equation has ΔH = −125 kJ, what is ΔH when the equation is reversed and all coefficients halved?
22. Calculate photon frequency for λ = 525 nm.
23. Calculate photon energy for ν = 6.00 × 10^14 s−1.
24. Write ground-state configurations for P, K+, and S2−.
25. Order Al3+, Mg2+, Na+, and F− by increasing radius.
26. Draw a valid Lewis structure for CH2O and identify geometry at carbon.
27. Draw nitrate resonance contributors and assign formal charges.
28. Predict geometry and polarity of PCl3.
29. Count sigma and pi bonds in CO2.
30. A gas occupies 2.20 L at 740 torr. Find volume at 1.00 atm and constant temperature.
31. Find moles in 10.0 L at 25.0 °C and 1.25 atm.
32. A mixture contains 0.20 mol N2 and 0.80 mol O2 at 1.00 atm total. Find partial pressures.
33. A gas is collected over water at total pressure 755 torr; water vapor pressure is 24 torr. Find dry gas pressure.
34. Rank expected boiling temperatures of CH4, NH3, and H2O qualitatively and justify.
35. A heating curve shows a plateau. Identify two measurements needed before assigning it as melting rather than boiling.

Answer key for quantitative items

1. 3.75 × 10^-4 m.
2. Volume 1.8 mL; density 8.7 g/mL to two significant figures.
3. (0.720)(62.930)+(0.280)(64.928)=63.489 u, approximately 63.49 u.
4. Ca3(PO4)2; (NH4)2SO3; Cu(NO3)2; N2O3.
5. Potassium dichromate; iron(III) oxide; tetraphosphorus decoxide; sodium hydrogen carbonate/bicarbonate.
6. 7.47 mol H2O.
7. 1.81 × 10^24 O atoms because each formula unit contains 12 O atoms.
8. Approximately 47.96% O using M(CaCO3) ≈ 100.09 g/mol.
9. For 100 g: approximately C 4.340, H 13.03, O 2.171 mol; divide by 2.171 → 2:6:1; C2H6O.
10. 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O.
11. 1.20 mol O2.
12. H2 is limiting; 2.40 mol NH3.
13. 83.3%.
14. 0.200 M approximately (5.85/58.44 = 0.1001 mol; divide by 0.5000 L).
15. 16.7 mL stock, diluted to 250.0 mL total.
16. Ca2+(aq) + CO3^2−(aq) → CaCO3(s).
17. 0.0175 mol Na+.
18. Moles base 0.002829 mol; 1:1; HCl 0.0943 M.
19. q=(80.0)(4.184)(22.5)=7.53 kJ.
20. ΔT = 1850/[(100.0)(4.184)] = 4.42 °C; solution warms and reaction heat is negative.
21. +62.5 kJ.
22. 5.71 × 10^14 s−1.
23. 3.98 × 10^-19 J per photon.
24. P: [Ne]3s2 3p3; K+: [Ar]; S2−: [Ar].
25. All are isoelectronic with ten electrons except check: Al3+, Mg2+, Na+, F− each has 10. Increasing radius: Al3+ < Mg2+ < Na+ < F−.
26. H2C=O with two lone pairs on O; carbon is trigonal planar, approximately 120°.
27. Three equivalent contributors; N commonly has formal charge +1, two singly bonded O atoms −1 each, double-bonded O 0, net −1.
28. Four electron domains; trigonal pyramidal molecular geometry; polar.
29. Two sigma and two pi bonds.
30. Convert 740 torr = 0.974 atm; V2=P1V1/P2=(0.974)(2.20)=2.14 L.
31. n=PV/RT=(1.25)(10.0)/[(0.082057)(298.15)] = 0.511 mol.
32. Mole fractions 0.20 and 0.80; partial pressures 0.20 atm N2, 0.80 atm O2.
33. 731 torr.
34. Expected order CH4 < NH3 < H2O; methane has dispersion only, ammonia hydrogen-bonds, and water forms an especially extensive hydrogen-bonding network. This is a qualitative comparison.
35. Pressure and the identity/composition of the sample, plus temperatures adjacent to the plateau; the phase assignment depends on the relevant phase-transition temperature under that pressure.

6. Scientific Explanation Rubric

Use for short answers, exam explanations, discussion, and capstone text.

7. Calculation Rubric

For a ten-point multi-step problem:

• 2 points: identifies knowns, wanted quantity, and applicable model.
• 2 points: correct symbolic setup or conversion path.
• 2 points: correct stoichiometric/equation relationship.
• 2 points: arithmetic and calculator execution.
• 1 point: units and dimensional consistency.
• 1 point: significant figures, scale, sign, and reasonableness.

Permit consequential credit when a later method is correct based on an earlier error.

8. Data-Story Capstone

Student brief

Choose one narrow, testable chemical claim connected to a real material or process. Examples:

• Why does road salt lower water's freezing behavior? Provide only a General Chemistry I preview if equilibrium treatment is not taught.
• How does molecular structure relate to the volatility of selected household compounds?
• How much carbon dioxide is predicted by combustion of a specified fuel quantity?
• How does battery-material composition affect theoretical mass contribution?
• What does a public air-quality or atmospheric composition dataset show at the chemical level?
• How does caloric energy compare with chemical enthalpy units?
• Which molecular features help explain a material's boiling point or solubility?

Required components

1. A precise claim.
2. At least two credible sources, including one authoritative data source.
3. One original calculation with units and assumptions.
4. One original data display or molecular/particle representation.
5. Explanation at macroscopic, particle, and symbolic levels.
6. Uncertainty, limitation, or boundary condition.
7. Accessible text alternatives and citations.

Rubric