Chemistry in the Kitchen and Beyond: Matter, Energy, and Change

This is where your child learns to describe stuff in the kitchen using numbers and evidence instead of just appearance — how heavy something is for its size (density), whether it's one substance or several mixed together, and whether a change is 'the same stuff rearranged' or 'now it's actually a different substance.' It also teaches that a measurement can only be as precise as the tool that made it — writing down more decimal places than your ruler can actually support is a false claim, not a nicety.

1. Significant figures as a claim about instrument precision+

   Given a measuring instrument (graduated cylinder, balance, ruler), record a measurement to the correct number of significant figures and identify the uncertainty implied by that instrument.

2. Significant figure rules in calculated (derived) quantities+

   Apply the rules for significant figures in addition/subtraction versus multiplication/division to compute a derived quantity (e.g., density) from two measured values, rounding the answer to the correct number of sig figs.

3. The element/compound/mixture/pure-substance classification scheme+

   Classify a given sample of matter as an element, compound, homogeneous mixture, or heterogeneous mixture by citing evidence of composition (uniformity, separability, particle representation) rather than appearance alone.

4. The distinction between physical and chemical change based on identity of substance+

   Distinguish physical from chemical change in a novel, previously unseen example by identifying whether the substance's chemical identity changed, not by matching the example to a memorized list of 'signs.'

5. The insufficiency of surface indicators as evidence for chemical change+

   Explain why color, texture, or the production of bubbles/gas/light/heat is not by itself sufficient evidence to classify a change as chemical, using the counter-cases of dissolving CO2 escaping soda (physical) and rusting (chemical, slow, no visible bubbles).

6. Density as an identifying physical property+

   Given density data (mass and volume, or mass and displaced volume) for an unknown material, calculate density and use it to identify the material from a reference table, justifying the identification as evidence rather than assumption.

7. Separation techniques matched to the physical property each exploits+

   Plan a two-step separation procedure (choosing from filtration, evaporation, and chromatography) for a novel kitchen mixture never used in class, justifying each step by the specific physical property it exploits (particle size, boiling point, solubility/polarity affinity).

8. Claim-evidence-reasoning structure applied to component identification via density+

   Construct a claim-evidence-reasoning argument that identifies each recovered component of a separated mixture, using measured density data as evidence and the definition of a pure substance as the reasoning link.

9. Instrument-limited precision as a general principle of quantitative claims+

   Generalize the principle that measurement precision must never exceed what the instrument can support to a context outside chemistry (e.g., a nutrition label, a news report of a poll percentage, a GPS coordinate) and explain why the reported precision is or is not justified.

Up to now the periodic table has just been a chart to look things up on. This unit turns it into a record of what's actually inside an atom — protons, neutrons, electrons, and how electrons are arranged — and shows how that arrangement explains patterns like which elements are metals, which react similarly, and why atomic size changes as you move around the table. It also covers the historical detective work (cathode rays, the gold foil experiment) that changed what scientists thought atoms looked like.

1. Subatomic particle counts and their relationship to atomic number, mass number, and charge+

   Given the number of protons, neutrons, and electrons in a neutral atom or ion, state the atomic number, mass number, and net charge.

2. Isotope notation and the neutron-count difference between isotopes of one element+

   Given the standard notation for an isotope (mass number and atomic number), determine the number of neutrons and distinguish it from other isotopes of the same element.

3. Electron configuration of main-group atoms+

   Write the electron configuration for a given main-group element in periods 1-4 using energy-level filling rules.

4. The historical revision of atomic models in response to specific experimental evidence (gold foil scattering)+

   Explain how the gold foil scattering evidence forced revision from the Thomson 'plum pudding' model to the Rutherford nuclear model.

5. The mechanism underlying periodic trends in atomic radius and ionization energy+

   Given unlabeled radius and ionization energy data for a set of elements not used in instruction, generate a rule connecting position on the periodic table to both trends and identify the underlying mechanism (nuclear charge and shielding).

6. Periodic trends (atomic radius, electronegativity, ionization energy) as a function of table position+

   Predict the relative atomic radius, electronegativity, and ionization energy of elements in an unfamiliar region of the periodic table (not drilled in class) and justify each prediction from position alone.

7. The relationship between atomic emission line spectra and quantized electron energy levels+

   Given an atomic emission spectrum (a set of discrete line positions) for an element, construct an argument for what the pattern of discrete lines implies about electron energy levels within the atom.

8. The Bohr model versus the electron cloud model as different-purpose simplifications+

   Compare the Bohr model and the electron cloud (quantum) model of the atom in terms of what each model can and cannot correctly predict.

9. Classification of elements by table position (metal/nonmetal/metalloid, group family)+

   Classify a given element as a metal, nonmetal, or metalloid, and as belonging to a specific group (alkali metal, halogen, noble gas, transition metal) based on its position on the periodic table.

10. The relationship between valence electron count and the charge of the ion an atom tends to form+

    Given the number of valence electrons for a main-group element, predict the ionic charge that element is most likely to form.

This is where the electron ideas from Unit 2 turn into an explanation for why atoms stick together at all. Your child will build ionic compounds (formula units) and covalent molecules (Lewis structures), figure out when a bond is polar versus nonpolar versus ionic, and learn to name and write chemical formulas as a skill, not a puzzle. A key move: comparing CO2 and H2O side by side to see that a molecule can have polar bonds but still be nonpolar overall, depending on shape.

1. Valence electron count and its relationship to group number and bonding tendency+

   Given an element's group number, state its number of valence electrons and whether it tends to lose, gain, or share electrons to reach a noble-gas configuration.

2. Charge-balancing procedure for binary ionic formula units+

   Construct the formula unit for a binary ionic compound from two given main-group elements by balancing positive and negative charge to net zero.

3. Lewis structure construction procedure for simple covalent molecules+

   Draw a correct Lewis structure for a simple covalent molecule (up to 3 atoms, single bonds only) by distributing valence electrons to satisfy the octet rule.

4. Electrostatic stability as the causal mechanism for electron transfer/sharing+

   Explain why an atom transfers or shares electrons in terms of relative electrostatic stability, without using goal-directed language ('wants', 'tries').

5. Electronegativity difference as the basis for bond-type classification along a continuum+

   Given electronegativity values or qualitative electronegativity trend position for two bonded atoms, classify the bond as nonpolar covalent, polar covalent, or ionic.

6. The distinction between individual bond polarity and overall molecular polarity as determined by shape/symmetry+

   Given contrasted molecular data (bond polarity vs. molecular symmetry) for two molecules such as CO2 and H2O, infer and justify whether each whole molecule is polar, distinguishing bond polarity from molecular polarity.

7. Polyatomic ion naming and formula-writing convention, including use of parentheses for multiple polyatomic units+

   Name and write the correct formula for an ionic compound containing a polyatomic ion, given the ion's name and charge from a reference table.

8. The causal link between bond/structure type and macroscopic properties (melting point, conductivity, solubility)+

   Given melting point, conductivity, and solubility data for an unfamiliar substance never discussed in class, argue in writing whether the substance is ionic, covalent-polar, or covalent-nonpolar, citing specific data points as evidence for the structural claim.

9. Textual evidence in scientific explanatory text signaling the limits of a simplified bonding model+

   Given a short excerpt from a chemistry reference text describing an exception to the octet rule (e.g., BF3 or SF6), determine what specific textual evidence signals the model's limitation and explain why the octet model fails in that case.

10. Relationship between ionic charge magnitude and lattice attraction strength (qualitative)+

    Predict the relative melting point ranking of two ionic compounds never discussed in class, using only their ions' charge magnitudes as evidence, and justify the prediction in terms of electrostatic attraction strength.

This is the quantitative heart of the whole course: how you count particles too small to see, by weighing them instead. Your child learns the mole as a fixed count (like 'a dozen,' but for a huge number), separate from mass, then learns to balance chemical equations as conservation of atoms, and finally uses balanced equations to predict how much product a reaction will make. It ends with a real sealed-bag reaction where they predict a result mathematically and then test it.

1. Molar mass of a compound calculated from atomic masses+

   Given a chemical formula, calculate its molar mass by summing atomic masses from the periodic table.

2. Mole-mass-particle conversions using molar mass and Avogadro's number+

   Convert a given mass of a pure substance to number of moles and to number of particles using molar mass and Avogadro's number.

3. The distinction between a counting unit and a mass or volume unit+

   Explain why the mole is defined as a fixed count of particles rather than a unit of mass or volume, using the dozen/golf-ball/bowling-ball comparison as evidence.

4. Balanced chemical equations with conserved atom counts+

   Balance a chemical equation by adjusting coefficients so that the number of atoms of each element is equal on both sides, without altering subscripts.

5. Mole-ratio stoichiometric conversion from mass of reactant to mass of product+

   Given a balanced equation and a known mass of one reactant, calculate the mass of a specified product using mole ratios and molar mass conversions.

6. Limiting reactant identified through ratio comparison+

   Identify the limiting reactant in a qualitative scenario (e.g., a recipe with a capped ingredient) by comparing the ratio of available reactants to the required ratio.

7. Conservation of mass applied to a novel open/closed reaction system+

   Apply conservation of atoms to explain an observed change in measured mass during a reaction in an unfamiliar closed or open system not used in instruction.

8. Full stoichiometric solution path applied to a novel reaction+

   Given an unpracticed reaction and a mass of reactant, generate the complete solution path (balance, convert to moles, apply ratio, convert to mass of product) without being told which steps apply.

9. Discrepancy between theoretical yield and measured lab result+

   Compare predicted (theoretical) mass of product to measured mass from the sealed-bag lab and construct an evidence-based explanation for any discrepancy.

If gases are just countless tiny particles bouncing around, why do they behave so predictably? This unit builds that particle picture (kinetic molecular theory) and uses it to explain pressure, and to derive the classic gas laws (Boyle's and Charles's) from real data before naming them. It ends at the ideal gas law, which uses the mole conversion skill from Unit 4 directly, and finishes with your child designing their own gas experiment (a balloon in hot and cold water, or a syringe) and explaining any mismatch between prediction and result.

1. The four assumptions of kinetic molecular theory+

   State the four core assumptions of kinetic molecular theory (particles in constant random motion, negligible particle volume relative to container, no intermolecular attraction, elastic collisions) and identify which assumption is violated in a given real-gas scenario.

2. The collision-based mechanism of gas pressure and its relationship to particle kinetic energy+

   Explain why gas pressure results from the frequency and force of particle collisions with container walls, connecting an increase in temperature to an increase in average particle kinetic energy and collision force.

3. The inverse P-V relationship and the direct V-T relationship in a fixed quantity of gas (Boyle's and Charles's Laws)+

   Given a pressure-volume or volume-temperature data set for a fixed amount of gas, classify the relationship as directly or inversely proportional and generalize it into a rule (Boyle's Law or Charles's Law) without being told the law's name in advance.

4. Algebraic solution of Boyle's, Charles's, and combined gas law equations+

   Solve for an unknown pressure, volume, or temperature using Boyle's Law, Charles's Law, or the combined gas law, correctly converting temperature to Kelvin before substitution.

5. Selection of the applicable gas law based on which variables are held constant or changing in a novel scenario+

   Given an unfamiliar gas scenario description (no equation or law named), determine which variable is held constant and select and apply the correct gas law (Boyle's, Charles's, combined, or ideal gas law).

6. Molar volume at STP and its relationship to moles, mass, and identity of a gas+

   Calculate moles, mass, or volume of a gas at STP using molar volume (22.4 L/mol) and the mole concept from Unit 4, distinguishing cases where mole count, mass, and volume diverge for different gases.

7. The ideal gas law equation PV=nRT combined with mole-mass conversion+

   Apply the ideal gas law (PV = nRT) to solve for an unknown quantity (P, V, n, or T) in a multi-step problem that first requires converting mass to moles using molar mass from Unit 4.

8. Conditions and mechanisms of real gas deviation from ideal gas behavior+

   Explain, using the kinetic molecular model, at least one specific condition (high pressure or low temperature) under which real gas behavior deviates measurably from ideal gas law predictions, citing particle volume or intermolecular attraction as the mechanism.

9. An original gas law prediction tested against experimental data from a kitchen-safe investigation+

   Design and carry out a kitchen-safe gas investigation (balloon in hot/cold water or syringe compression), collecting quantitative data and using it to test a specific gas law prediction, then reconciling any discrepancy between predicted and observed results using KMT rather than attributing it solely to measurement error.

10. The distinction between mechanistic (particle-collision) and goal-directed/substance-transfer explanations of gas behavior+

    Compare and critique a peer's or provided sample explanation of a gas behavior scenario, identifying whether the explanation relies on goal-directed or substance-transfer language (e.g., 'the gas wants to expand,' 'heat enters the balloon') versus correct particle-collision mechanism.

This unit answers why things dissolve (using the polarity ideas from Unit 3) and how to measure how much is dissolved (using the mole math from Unit 4). Your child will figure out 'like dissolves like' by testing solutes in water versus oil, then move into molarity and dilution calculations, and finally build and read a solubility curve from real data they collect by heating water and dissolving a solid in it to the saturation point.

1. The relationship between molecular polarity/shape and solubility ('like dissolves like')+

   Given a solute and solvent pair not used in class demonstrations, students will classify whether it will dissolve using bond polarity and molecular shape reasoning.

2. The molarity formula and its use given moles or mass of solute and solution volume+

   Students will calculate the molarity of a solution given moles or grams of solute and volume of solution, using M = mol/L.

3. The dilution formula M1V1 = M2V2 and solving for an unknown quantity+

   Students will calculate a new molarity or volume after dilution using M1V1 = M2V2, given three of the four quantities.

4. Conservation of moles of solute during dilution+

   Students will explain why the total moles of solute remain constant during dilution even though the solution appears more dilute, in response to a claim that dilution destroys some of the solute.

5. Sequencing of mole conversion, molarity, and dilution formulas in an unfamiliar multi-step problem+

   Given a novel multi-step problem combining mole conversion, molarity, and dilution with numbers not matching any practiced problem, students will select and execute the correct sequence of formulas.

6. Dynamic equilibrium in a saturated solution+

   Students will explain that a saturated solution at equilibrium continues to have particles dissolving and recrystallizing at equal rates, rather than describing dissolving as having stopped.

7. Factors affecting dissolving rate explained through particle collision frequency+

   Students will predict the effect of temperature, agitation, and surface area on the RATE of dissolving, citing particle collision frequency from the kinetic-molecular model, for a scenario not explicitly demonstrated in class.

8. The solubility curve as a graphical representation of saturation point versus temperature+

   Students will construct a solubility curve graph from experimentally collected temperature and mass-dissolved-at-saturation data, with correctly labeled axes and units.

9. Classification of solution states (unsaturated, saturated, supersaturated) relative to a solubility curve+

   Given an unfamiliar solubility curve for a solute not studied in class, students will classify a data point as unsaturated, saturated, or supersaturated and justify the classification using the curve's shape.

10. The limits of the 'like dissolves like' model for molecules with mixed polar and nonpolar regions+

    Given a completely unfamiliar polar molecule with an unusual structure (e.g., an amino acid or surfactant not discussed in class) and asked whether 'like dissolves like' correctly predicts its solubility behavior, students will identify the specific structural feature that makes the simple rule an incomplete explanation and generalize a refined account.

This unit shifts the question from 'how much' to 'how fast, and does it ever actually stop.' Your child explains reaction speed using particle collisions (frequency and energy), and equilibrium as a state where forward and reverse reactions are happening at equal rates — not a state where reaction has stopped. It moves through rate factors and energy diagrams (fully explained with worked examples), then into equilibrium and Le Chatelier's principle (mostly figured out by observing a real reversible system), and ends with Keq expressions and a rate-investigation lab.

1. The four factors affecting reaction rate (concentration, temperature, surface area, catalyst)+

   Given a reaction scenario, state which of the four rate factors (concentration, temperature, surface area, catalyst) was changed and predict whether rate increases or decreases.

2. Collision theory as the mechanism linking temperature to reaction rate+

   Explain, using a collision-frequency argument, why increasing temperature increases reaction rate through two distinct mechanisms: more frequent collisions and a greater fraction of particles with energy above the activation energy threshold.

3. Energy diagrams showing activation energy and overall energy change+

   Interpret an unfamiliar energy diagram (not used in instruction) to identify activation energy, whether the reaction is exothermic or endothermic, and the relative energy of reactants versus products.

4. Dynamic equilibrium as a state of equal forward and reverse rates+

   Given observational data from a reversible color-change or dissolving/precipitating system over time, classify whether the system has reached dynamic equilibrium or is still net-shifting toward one side.

5. The distinction between dynamic equilibrium and a static or stopped state+

   Explain why a system at equilibrium is still undergoing reaction in both directions, using the rate-equality definition rather than describing it as stopped or balanced 50/50.

6. Le Chatelier's principle applied to concentration and temperature stressors+

   Predict the direction an equilibrium system will shift when a familiar type of stressor (concentration change or temperature change, both practiced in class) is applied, and justify the prediction using rate reasoning.

7. Le Chatelier's principle extended to an unpracticed stressor type+

   Predict and justify the equilibrium shift for a stressor type not explicitly practiced in class (e.g., a volume/pressure change in a gas-phase system, or continuous removal of a product), using the same rate-based reasoning developed for familiar stressors.

8. The equilibrium constant expression and its relationship to balanced-equation mole ratios+

   Write a correct equilibrium constant (Keq) expression for a given balanced chemical equation, using the coefficients as exponents.

9. The relationship between the magnitude of Keq and the relative amounts of products versus reactants at equilibrium+

   Given a Keq expression and its numeric value for an unfamiliar reaction, infer whether the reaction favors products or reactants at equilibrium and compare two such reactions' relative extent.

10. A controlled experimental design isolating one rate factor using tablet-dissolving reaction time as the dependent measure+

    Plan and carry out a controlled investigation varying one factor (surface area or temperature) affecting the rate of a tablet-dissolving reaction, collecting time-based rate data with appropriate controls.

11. Experimental rate data from the tablet-dissolving investigation+

    Analyze a peer or provided data set from a rate investigation (time vs. amount reacted, under varied surface area/temperature) to determine which condition produced the fastest rate and identify a plausible source of experimental error.

The year's final unit ties everything together: acid-base reactions and redox (electron-transfer) reactions are both explained as one particle moving from one substance to another — a proton in acid-base chemistry, an electron in redox. Your child moves from the simple Arrhenius definition to the more general Bronsted-Lowry one, tests household substances with homemade red cabbage indicator, runs a real titration to find an unknown concentration, and reframes rusting, combustion, and batteries as electron-transfer reactions.

1. Arrhenius vs. Bronsted-Lowry acid-base definitions+

   State the Arrhenius and Bronsted-Lowry definitions of acids and bases and identify which one applies to a given reaction.

2. Conjugate acid-base pairs in proton-transfer reactions+

   Identify the conjugate acid-base pair in a Bronsted-Lowry proton-transfer reaction, including in reactions not shown during instruction.

3. PH scale as a measure of relative H+ concentration, and indicator color response+

   Predict the relative pH and color change of a red cabbage indicator solution for common household substances based on their particle-level proton concentration.

4. Titration stoichiometry (moles, volume, molarity relationships)+

   Calculate the unknown molar concentration of an acid or base from titration volume and concentration data using mole ratios from a balanced neutralization equation.

5. Titration stoichiometry with non-1:1 mole ratios+

   Determine an unknown solution's concentration from titration data when the acid or base is polyprotic or the mole ratio is not 1:1.

6. Oxidation number assignment rules+

   Assign oxidation numbers to atoms in a compound or ion using periodic-trend-based rules.

7. Oxidation and reduction identification via oxidation-number change+

   Identify the species oxidized and the species reduced in an unfamiliar redox reaction by tracking change in oxidation number.

8. Redox reactions in everyday contexts (rusting, combustion, batteries)+

   Explain rusting, combustion, and battery discharge as everyday examples of electron transfer between specific chemical species.

9. Argument from pH data to particle-level proton-transfer claims+

   Construct a written, evidence-based argument connecting measured pH data to a particle-level claim about relative proton concentration and direction of proton transfer.

10. Titration lab design and execution+

    Design and carry out a titration procedure to determine the unknown concentration of a household acid or base, selecting appropriate indicator and calculation method.

11. Structural analogy between proton transfer and electron transfer reactions+

    Compare the particle transferred (proton vs. electron) across acid-base and redox reactions and explain why both are classified as transfer reactions.