Midterm Study Guide for Chemistry PDF

Significant Figures (Sig Figs) * Counting Sig Figs: * Non-zero digits are always significant. * Any zeros between significant digits are significant. * Leading zeros (zeros to the left of the first non-zero digit) are not significant. * Trailing zeros are significant if there is a...

Significant Figures (Sig Figs) * Counting Sig Figs: * Non-zero digits are always significant. * Any zeros between significant digits are significant. * Leading zeros (zeros to the left of the first non-zero digit) are not significant. * Trailing zeros are significant if there is a decimal point. * Math with Sig Figs: * Addition/Subtraction: Round the answer to the least number of decimal places. * Multiplication/Division: Round the answer to the least number of significant figures. ________________ 2. Unit Conversions * mL → L: Divide by 1000 (1 L = 1000 mL) * L → mL: Multiply by 1000 * g → mg: Multiply by 1000 (1 g = 1000 mg) * mg → g: Divide by 1000 * km → m → cm → mm: * 1 km = 1000 m * 1 m = 100 cm * 1 cm = 10 mm ________________ 3. Atoms, Elements, and Compounds * Atom: The smallest unit of an element that retains its properties. * Element: A pure substance made up of only one type of atom. * Compound: A substance made up of two or more different elements chemically bonded. ________________ 4. States of Matter * Solid: Definite shape and volume. Particles are closely packed. * Liquid: Definite volume, but takes the shape of its container. Particles can move around but are still close. * Gas: No definite shape or volume. Particles are far apart and move freely. ________________ 5. Physical vs. Chemical Change * Physical Change: Changes that do not alter the chemical composition (e.g., melting, freezing, cutting). * Chemical Change: A change that results in the formation of new chemical substances (e.g., combustion, rusting). ________________ 6. Physical vs. Chemical Properties * Physical Property: Properties that can be observed or measured without changing the substance (e.g., color, melting point, density). * Chemical Property: Describes a substance’s ability to undergo a chemical change (e.g., flammability, reactivity with acids). ________________ 7. Dalton's Atomic Theory * All matter is made up of atoms. * Atoms of the same element are identical. * Atoms combine in simple whole-number ratios to form compounds. * Chemical reactions involve the rearrangement of atoms. ________________ 8. Atomic Models * J.J. Thomson: Discovered the electron using the cathode ray tube; * Ernest Rutherford: Conducted the gold foil experiment; concluded the atom has a dense, positively charged nucleus and is mostly empty space. * Niels Bohr: Proposed that electrons orbit the nucleus in fixed paths (orbits), and energy levels are quantized. ________________ 9. Structure of the Atom * Protons: Positive charge, located in the nucleus. * Neutrons: Neutral charge, located in the nucleus. * Electrons: Negative charge, located in the electron cloud outside the nucleus. ________________ 10. Periodic Table * Periodic: Refers to repeating trends or patterns. * Elements are arranged by increasing atomic number. * Groups/Families: Vertical columns; elements in the same group have similar chemical properties. * Periods: Horizontal rows; properties change gradually across a period. ________________ 11. Families on the Periodic Table * Alkali Metals: Group 1 (Li, Na, K), highly reactive, especially with water. * Alkaline Earth Metals: Group 2 (Be, Mg, Ca), reactive, but less so than alkali metals. * Transition Metals: Groups 3-12, can form multiple charges. * Halogens: Group 17 (F, Cl, Br), highly reactive, form salts with metals. * Noble Gases: Group 18 (He, Ne, Ar), inert, non-reactive. ________________ 12. Isotopes * Isotopes are atoms of the same element that have different numbers of neutrons and therefore different masses. ________________ 13. Mass Spectrum * A mass spectrum shows the relative abundance of isotopes of an element. You can determine the element’s identity based on its mass spectrum. ________________ 14. Naming Ionic Compounds * Binary Ionic Compounds: Composed of two elements (e.g., NaCl, CaO). Name the cation (metal) first, then the anion (non-metal). * Ternary Ionic Compounds: Contain a polyatomic ion (e.g., NaNO₃, CaCO₃). Name the cation first, then the polyatomic ion. * Ionic Compounds with Multiple Charges: Transition metals often have multiple charges (e.g., Fe²⁺, Fe³⁺). Indicate the charge in Roman numerals (e.g., Iron(III) chloride). ________________ 15. Empirical Formula * The empirical formula is the simplest ratio of elements in a compound. Assume you have 100 grams of that sample and then using the percentages in the question find the percent of each element. Then divide each element by the molar mass.Then divide all the mole amounts by the smallest mole amount and then you will get whole number ratios. ________________ 16. Percent Composition * Percent composition = (mass of element in compound / molar mass of compound) × 100. ________________ 17. Laws of Conservation in Chemical Reactions * The Law of Conservation of Mass states that mass is neither created nor destroyed in a chemical reaction, so when balancing a chemical equation, the number of atoms on each side must be equal. ________________ 18. Kinetic Molecular Theory of Gases * Gases are made of tiny particles in constant motion. * Gas particles are far apart, have negligible volume, and experience elastic collisions. * The average kinetic energy of gas particles is directly proportional to temperature. ________________ 19. Gas Laws * Boyle’s Law: P1V1=P2V2 2​ (Pressure and volume are inversely related at constant temperature). * Charles’s Law: V1T1=V2T2 (Volume and temperature are directly related at constant pressure). * Gay-Lussac’s Law: P1T1=P2T2 (Pressure and temperature are directly related at constant volume). * Ideal Gas Law: PV=nRTRT (Relates pressure, volume, temperature, and moles of gas). * Dalton’s Law of Partial Pressures: PTotal=P1+P2+P3+… * Graham’s Law: Rate of effusion m1 /Rate of effusion m2=M2/M1(The rate of effusion of a gas is inversely proportional to the square root of its molar mass). ________________ 20. Stoichiometry * Mole-to-Mole Conversions: Use the balanced chemical equation to convert between moles of reactants and products. * Limiting and Excess Reactants: * The limiting reactant is the reactant that runs out first and determines the amount of product formed. * The excess reactant is the reactant that is left over after the reaction is complete. * Percent Yield: Percent Yield=Actual Yield - Theoretical Yield×100 ________________ 21. Sample Problem Types * Empirical Formula Example: If you have 40.0 g of carbon and 60.0 g of oxygen, the empirical formula of the compound is CO₂. * Mole-to-Mole Conversions: Given 2 moles of H₂O, how many moles of O₂ are needed? Use the balanced equation to set up the conversion. * Limiting Reactant Problem: If you have 5.00 g of Na and 10.00 g of Cl₂, which is the limiting reactant? Use molar masses to convert to moles and compare mole ratios.