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AP Chem Unit 2
AP Chemistry
| Term | Definition | Dive Deep | Memory Hack |
|---|---|---|---|
| patterns of chemical reactivity | Recurring ways chemicals react and combine. | Understanding patterns helps predict products—like how metals tend to lose electrons and nonmetals gain them. | Chemistry has rhythm: learn the patterns, predict the dance. |
| balancing equations | Making sure both sides of a chemical equation have the same number of each atom. | Matter cannot be created or destroyed, so every atom that goes in must come out. | Law of conservation’s accountant. |
| law of conservation of mass | Mass is neither created nor destroyed in chemical reactions. | The total mass of reactants equals the total mass of products. | What goes in must come out. |
| synthesis reaction | A reaction where two or more substances combine to form one product. | General form: A + B → AB. Example: 2H₂ + O₂ → 2H₂O. | Putting pieces together. |
| decomposition reaction | A reaction where one compound breaks down into two or more simpler substances. | General form: AB → A + B. Example: 2H₂O → 2H₂ + O₂. | Breaking things apart. |
| combustion reaction | A reaction where a substance reacts with oxygen to produce CO₂ and H₂O. | Usually involves hydrocarbons and releases energy as heat and light. | Fire reaction: fuel + O₂ → CO₂ + H₂O. |
| writing neutral compounds | Creating chemical formulas so overall charge equals zero. | Balance positive and negative charges using subscripts. Example: Na⁺ and Cl⁻ form NaCl. | Charge balance = chemical harmony. |
| formula weight | The sum of the atomic weights of atoms in a formula. | Measured in amu (atomic mass units). Used for molecular or ionic compounds. | The molecule’s weight tag. |
| Avogadro’s number | The number of particles in one mole of a substance (6.022 × 10²³). | Applies to atoms, ions, or molecules; connects microscopic and macroscopic worlds. | The chemist’s counting bridge. |
| the mole | A counting unit representing 6.022 × 10²³ particles. | Links atoms to grams, making chemical quantities measurable in labs. | The chemist’s dozen. |
| why we use moles | To convert between atoms/molecules and grams in real-world amounts. | Atoms are too small to count directly, so moles scale them up. | The zoom tool for chemistry. |
| mole triangle | A diagram showing the relationships between moles, mass, and number of particles. | Top of the triangle: mass ↔ moles ↔ number of particles. | Use the triangle to jump between grams, moles, and atoms. |
| conversions | Changing between units (grams, moles, particles, liters). | Use conversion factors from molar mass or Avogadro’s number. | Dimensional analysis = unit gymnastics. |
| percent composition | The percentage by mass of each element in a compound. | Calculated by (mass of element ÷ total mass) × 100. | Find the recipe of a compound. |
| empirical formula from % composition | The simplest ratio of elements in a compound derived from percentages. | Convert % to grams, grams to moles, then divide by smallest. | Simplify to lowest whole-number ratio. |
| molecular formula from empirical formula | The actual formula based on molar mass. | Multiply the empirical formula by n, where n = (molar mass ÷ empirical mass). | Scale up the simplest version. |
| how EF fits into MF | Shows how many times the empirical formula repeats in the molecular formula. | EF × n = MF, where n is an integer. | E.F. is the building block; M.F. is the full model. |
| combustion analysis | A lab method to determine empirical formula by burning a compound and measuring CO₂ and H₂O produced. | Used for organic compounds to find C, H, and O ratios. | Burn it to reveal its recipe. |
| mole-to-mole relationship | Stoichiometric ratios between substances in a balanced equation. | Coefficients show how many moles of each react. | Recipe math for reactions. |
| percent yield | (Actual yield ÷ Theoretical yield) × 100. | Measures efficiency of a reaction. Some product is always lost. | Chemistry’s report card. |
| limiting reactant | The reactant that runs out first, stopping the reaction. | Determined by comparing mole ratios of reactants to the balanced equation. | The bottleneck of the reaction. |
| formula of an unknown hydrate | Finding the ratio of water to salt in a hydrate compound. | Heat to drive off water, find moles of each, form ratio. | Dehydrate to discover its formula. |
| chemical changes and equations | Changes that produce new substances, shown in chemical equations. | Signs: color change, gas formation, heat, light, precipitate. | When atoms rearrange, chemistry happens. |
| precipitation reaction | A reaction forming an insoluble solid (precipitate) when two aqueous solutions mix. | Use solubility rules to predict the solid. | Solid surprise reaction. |
| solubility rules | Rules predicting if ionic compounds dissolve in water. | Example: Nitrates (NO₃⁻) always soluble; most Ag⁺, Pb²⁺, Hg₂²⁺ salts are insoluble. | The guidebook for dissolving. |
| acid-base neutralization reaction | An acid reacts with a base to form water and a salt. | Example: HCl + NaOH → NaCl + H₂O. | Acid + base = neutral taste. |
| gas formation reaction | A reaction producing a gas such as CO₂, H₂, or NH₃. | Often occurs in acid–carbonate or redox reactions. | Bubbles mean new gas formed. |
| oxidation and reduction reactions | Reactions where electrons are transferred between species. | Oxidation = loss of electrons; reduction = gain. (OIL RIG). | Who lost or gained electrons? |
| electrolytes | Substances that dissociate into ions in water and conduct electricity. | Strong electrolytes dissociate completely, weak ones partially. | Charged swimmers in solution. |
| writing net ionic equations | Equations showing only the ions that change in a reaction. | Remove spectator ions; balance charge and atoms. | The essential story of the reaction. |
| solutions | Homogeneous mixtures of solute dissolved in solvent. | Properties depend on solute concentration and interactions. | The uniform mix. |
| molarity | Concentration measured as moles of solute per liter of solution. | M = moles ÷ liters. | Concentration shorthand. |
| m = moles of solute / volume of solution in liters | The formula for molarity. | Used to prepare and calculate concentrations. | M = n/V. |
| dilutions | Reducing solution concentration by adding solvent. | Use M₁V₁ = M₂V₂ to find new concentrations. | Add water, same moles, more volume. |
| how something dissolves (solvation ionic) | Ions separate and are surrounded by water molecules. | The polar water stabilizes ions via hydration shells. | Water hugs the ions apart. |
| ideal gas at STP | At 1 atm and 273 K, one mole of gas occupies 22.4 L. | Used for gas laws and stoichiometry with gases. | Standard gas yardstick. |
Created by:
Leo12345
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