1. - Chemistry Topics Semester 1
1.1. - Bonds
1.1.1. - Ionic Bonds
1.1.1.1. - Formed between metals and nonmetals.
1.1.1.2. - Electrons are transferred from one atom to another, creating charged ions (e.g., NaCl, MgO).
1.1.1.3. - Properties: High melting/boiling points, soluble in water, conduct electricity when dissolved.
1.1.2. - Covalent Bonds
1.1.2.1. - Formed between nonmetals.
1.1.2.2. - Electrons are shared between atoms (e.g., H2O, CO2).
1.1.2.3. - Properties: Lower melting/boiling points, poor electrical conductors.
1.1.3. - Metallic Bonds
1.1.3.1. - Found in metals.
1.1.3.2. - "Sea of electrons" around positive metal ions allows for conductivity and malleability (e.g., Cu, Al).
1.1.4. - Lewis and Lattice Structures
1.1.4.1. - Lewis structures: Represent valence electrons in molecules.
1.1.4.2. - Lattice structures: 3D arrangement of ions in ionic compounds.
1.2. - Elements, Compounds, and Mixtures
1.2.1. - Elements: Pure substances made of only one type of atom (e.g., O2, He).
1.2.2. - Compounds: Substances made of two or more elements chemically bonded (e.g., H2O, NaCl).
1.2.3. - Mixtures: Two or more substances physically combined, not chemically (e.g., air, salad).
1.2.3.1. - Homogeneous mixtures (solutions): Uniform composition (e.g., saltwater).
1.2.3.2. - Heterogeneous mixtures: Non-uniform composition (e.g., oil and water).
1.3. - Law of Conservation of Mass
1.3.1. - Matter cannot be created or destroyed in a chemical reaction.
1.3.2. - Total mass of reactants = Total mass of products.
1.3.3. - Example: 2H2 + O2 → 2H2O (Mass of hydrogen + oxygen = Mass of water).
1.4. - Reactants and Products
1.4.1. - Reactants: Substances that undergo a chemical change.
1.4.2. - Products: Substances formed after the reaction.
1.4.3. - Example: In combustion of methane (CH4 + 2O2 → CO2 + 2H2O), CH4 and O2 are reactants; CO2 and H2O are products.
1.5. - Isotopes
1.5.1. - Atoms of the same element with different numbers of neutrons (e.g., Carbon-12 and Carbon-14).
1.5.2. - Relative Atomic Mass: Weighted average of isotopes based on abundance.
1.5.3. - Radioactivity: Unstable isotopes decay, releasing energy (e.g., Uranium-238).
1.6. - Chemistry Materials and Instruments
1.6.1. - Bunsen Burner
1.6.1.1. - Used to heat substances in labs.
1.6.1.2. - Flame can be adjusted by regulating air supply.
1.6.2. - Lab Skills and Safety
1.6.2.1. - Always wear safety goggles and gloves.
1.6.2.2. - Follow proper procedures for handling chemicals.
1.6.2.3. - Use fume hoods when working with volatile substances.
1.7. - Balancing Equations
1.7.1. - Ensure the number of atoms of each element is the same on both sides of the equation.
1.7.2. - Example: Unbalanced → H2 + O2 → H2O; Balanced → 2H2 + O2 → 2H2O.
1.8. - Periodic Table
1.8.1. - Metals
1.8.1.1. - Shiny, conductive, malleable, and ductile (e.g., Fe, Al).
1.8.2. - Nonmetals
1.8.2.1. - Poor conductors, brittle, can be gases or solids (e.g., O2, S).
1.8.3. - Metalloids
1.8.3.1. - Properties of both metals and nonmetals (e.g., Si, B).
1.9. - Element Families
1.9.1. - Alkali Metals: Highly reactive, 1 valence electron (e.g., Na, K).
1.9.2. - Halogens: Reactive nonmetals, 7 valence electrons (e.g., F, Cl).
1.9.3. - Noble Gases: Inert, full valence shells (e.g., He, Ne).
1.10. - Atomic Properties
1.10.1. - Ions
1.10.1.1. - Charged atoms (e.g., Na+, Cl-).
1.10.2. - Valence Electrons
1.10.2.1. - Electrons in the outermost shell.
1.10.3. - Electronegativity
1.10.3.1. - Tendency to attract electrons (e.g., F has the highest).
1.10.4. - Electron Affinity
1.10.4.1. - Energy change when an atom gains an electron.
1.10.5. - Ionization Energy
1.10.5.1. - Energy needed to remove an electron.
1.10.6. - Atomic Radius
1.10.6.1. - Size of an atom (decreases across a period, increases down a group).
1.10.7. - Period Trends
1.10.7.1. - Properties like electronegativity, atomic radius, and ionization energy vary predictably across periods and groups.
1.11. - Risk Assessment
1.11.1. - Identify potential hazards (e.g., chemical spills, burns).
1.11.2. - Take precautionary measures (e.g., wear PPE, proper labeling of chemicals).
1.12. - Types of Reactions
1.12.1. - Synthesis: A + B → AB (e.g., 2H2 + O2 → 2H2O).
1.12.2. - Decomposition: AB → A + B (e.g., 2H2O → 2H2 + O2).
1.12.3. - Single Replacement: A + BC → AC + B (e.g., Zn + 2HCl → ZnCl2 + H2).
1.12.4. - Double Replacement: AB + CD → AD + CB (e.g., NaCl + AgNO3 → NaNO3 + AgCl).
1.12.5. - Combustion: Hydrocarbon + O2 → CO2 + H2O (e.g., CH4 + 2O2 → CO2 + 2H2O).
1.13. - Physical and Chemical Changes and Properties
1.13.1. - Physical Change: Does not alter the substance's identity (e.g., melting ice, dissolving salt).
1.13.2. - Chemical Change: Produces a new substance (e.g., rusting iron, burning wood).
1.13.3. - Physical Properties: Observable without changing composition (e.g., color, density).
1.13.4. - Chemical Properties: Observed during chemical changes (e.g., reactivity, flammability).
2. Chemistry Topics Semester 2
2.1. Dimensional Analysis
2.1.1. Definition: A method to convert between units using conversion factors.
2.1.2. Key Steps:
2.1.2.1. Identify the given value and unit.
2.1.2.2. Determine the conversion factor(s).
2.1.2.3. Multiply by conversion factors to cancel units.
2.1.2.4. Ensure the final answer has the desired unit.
2.1.3. Examples:
2.1.3.1. Converting grams to moles.
2.1.3.2. Converting meters to kilometers.
2.1.4. Common Conversion Factors:
2.1.4.1. 1 inch = 2.54 cm
2.1.4.2. 1 mole = 6.022 × 10²³ particles
2.1.5. Definition: The SI unit for the amount of substance.
2.1.6. Avogadro’s Number: 6.022 × 10²³ particles/mol.
2.1.7. Relationships:
2.1.7.1. 1 mole = molar mass in grams.
2.1.7.2. 1 mole = 22.4 L (at STP for gases).
2.1.7.3. 1 mole = 6.022 × 10²³ atoms/molecules/ions.
2.1.8. Conversions:
2.1.8.1. Mass ↔ Moles: Moles = Mass / Molar Mass.
2.1.8.2. Particles ↔ Moles: Moles = Particles / Avogadro’s Number.
2.1.8.3. Volume ↔ Moles (for gases): Moles = Volume / 22.4 L.
2.1.9. Definition: The limiting reagent is the reactant that runs out first and determines the maximum amount of product formed.
2.1.10. Steps to Identify:
2.1.10.1. Convert reactants to moles.
2.1.10.2. Use stoichiometry to determine which reactant produces the least product.
2.1.10.3. The reactant that produces the least product is the **limiting reagent**.
2.1.11. Excess Reagent:
2.1.11.1. The reactant left over after the reaction.
2.1.11.2. Can be calculated by subtracting the used amount from the given amount.
2.1.12. Example Problem:
2.1.12.1. Given 10g of A reacts with 15g of B, determine the limiting reagent and the amount of product formed.
2.1.13. Definition: The calculation of reactants and products in a chemical reaction.
2.1.14. Mole Ratio: Derived from balanced chemical equations.
2.1.15. Steps:
2.1.15.1. Convert given quantities to moles.
2.1.15.2. Use mole ratio to find moles of desired substance.
2.1.15.3. Convert back to grams or other units if needed.
2.1.16. Example:
2.1.16.1. How many grams of CO₂ are produced from 10g of CH₄?
2.1.17. Used in real-world scenarios like:
2.1.17.1. Industrial chemical production.
2.1.17.2. Pharmaceutical drug formulation.
2.1.17.3. Environmental chemistry for pollutant analysis.
2.1.18. Moles
2.2. Moles
2.2.1. Definition: The SI unit for the amount of substance.
2.2.2. Avogadro’s Number: 6.022 × 10²³ particles/mol.
2.2.3. Relationships:
2.2.3.1. 1 mole = molar mass in grams.
2.2.3.2. 1 mole = 22.4 L (at STP for gases).
2.2.3.3. 1 mole = 6.022 × 10²³ atoms/molecules/ions.
2.2.4. Conversions:
2.2.4.1. Mass ↔ Moles: Moles = Mass / Molar Mass.
2.2.4.2. Particles ↔ Moles: Moles = Particles / Avogadro’s Number.
2.2.4.3. Volume ↔ Moles (for gases): Moles = Volume / 22.4 L.
2.3. Limiting and Excess Reagents
2.3.1. Definition: The limiting reagent is the reactant that runs out first and determines the maximum amount of product formed.
2.3.2. Steps to Identify:
2.3.2.1. Convert reactants to moles.
2.3.2.2. Use stoichiometry to determine which reactant produces the least product.
2.3.2.3. The reactant that produces the least product is the **limiting reagent**.
2.3.3. Excess Reagent:
2.3.3.1. The reactant left over after the reaction.
2.3.3.2. Can be calculated by subtracting the used amount from the given amount.
2.3.4. Example Problem:
2.3.4.1. Given 10g of A reacts with 15g of B, determine the limiting reagent and the amount of product formed.
2.4. Stoichiometry
2.4.1. Definition: The calculation of reactants and products in a chemical reaction.
2.4.2. Mole Ratio: Derived from balanced chemical equations.
2.4.3. Steps:
2.4.3.1. Convert given quantities to moles.
2.4.3.2. Use mole ratio to find moles of desired substance.
2.4.3.3. Convert back to grams or other units if needed.
2.4.4. Emperical Formula and Molecular Formula
2.4.4.1. First, you have to get the molar mass of the molecular formula
2.4.4.2. You are provided either the percentage or the grams out of 100 grams.
2.4.4.3. You then divide the mass by its molecular mass which you find in the periodic table to find the moles.
2.4.4.4. After that you divide the all the values by the smallest mole value.
2.4.4.5. Then you get numbers and you put them beside the chemical and if one has a decimal, you multiply or divide and then round the decimal number to the nearest number.
2.4.4.6. Then you have found the emperical formula
2.4.4.7. Find the molar mass of the emperical formula
2.4.4.8. Divide the molar mass of the molecular formula by the molar mass of the emperical formula.