Comprehensive High School Chemistry Final Exam Review Study Notes

Laboratory Safety and Instrumentation

• Lab Safety Protocols

  • 1. Always wear safety goggles to protect eyes from chemical splashes or debris.

  • 2. Tie back long hair and avoid loose clothing to prevent contact with flames or chemicals.

  • 3. Never eat, drink, or chew gum in the laboratory to avoid accidental ingestion of toxic substances.

• Identification and Use of Lab Instruments

  • Beaker: Used for stirring, mixing, and heating liquids; provides rough estimations of volume.

  • Graduated cylinder: Used for measuring precise volumes of liquids.

  • Triple-Beam balance: Used to measure the mass of an object in grams ($g$).

  • Eudiometer: A laboratory device that measures the change in volume of a gas mixture following a physical or chemical change.

  • Burette: Used in titrations to dispense known volumes of a liquid with high precision.

  • Test tube clamp: Used to hold a single test tube, particularly when heating or moving it.

  • Test tube tongs: Used specifically for lifting or carrying test tubes.

  • Crucible tongs: Designed to handle hot crucibles or other hot porcelain ware.

  • Beaker tongs: Used to safely handle and move hot beakers.

  • Iron ring clamp: Attaches to a ring stand to support laboratory glassware such as a beaker or funnel.

• Measurement Techniques

  • Recording Measured Values: Always record every certain digit from the markings on the instrument plus one final estimated digit (the digit of uncertainty).

  • Volume Comparison: If measuring $5\,mL$ of water:

    • Using a $10\,mL$ graduated cylinder, the measurement would be more precise (e.g., recorded as $5.00\,mL$).

    • Using a $100\,mL$ graduated cylinder, the measurement would be less precise (e.g., recorded as $5.0\,mL$) because the markings are further apart.

Properties of Matter and Physical States

• Matter: Defined as anything that has mass and takes up space (volume).

• Mass vs. Weight:

  • Mass: The amount of matter in an object; it remains constant regardless of location.

  • Weight: The force of gravity acting on an object's mass; it changes based on the gravitational pull of the location.

• Measuring Volume:

  • Regular Solid: Calculated using geometric formulas (e.g., $Length \times Width \times Height$).

  • Irregular Solid: Measured using the water displacement method (Final Volume - Initial Volume).

• Phases of Matter Notation:

  • s: Solid

  • l: Liquid

  • g: Gas

  • aq: Aqueous (dissolved in water)

• Elements vs. Compounds:

  • Element: A pure substance consisting of only one type of atom. Example: Gold ($Au$).

  • Compound: A substance formed when two or more chemical elements are chemically bonded together. Example: Water ($H_2O$).

• Density:

  • Definition: The ratio of mass to volume ($D = \frac{m}{v}$).

  • Density of Water: roughly $1.0\,g/cm^3$ or $1.0\,g/mL$.

  • Floating and Sinking: Objects with a density less than $1.0\,g/cm^3$ float in water; those with a density greater than $1.0\,g/cm^3$ sink.

• Table: Characteristics of States of Matter

Significant Figures and Mathematical Rules

• Identifying Significant Figures:

  • $100.0$: 4 significant figures (trailing zeros with a decimal point are significant).

  • $150$: 2 significant figures (trailing zeros without a decimal are placeholders).

  • $0.004$: 1 significant figure (leading zeros are never significant).

  • $0.00903801$: 6 significant figures.

  • $12$: 2 significant figures.

  • $1.0$: 2 significant figures.

• Addition and Subtraction Rule: The result must be rounded to the same number of decimal places as the measurement with the fewest decimal places.

  • a. $238.0 + 12.008 = 250.0$ (rounded to tenths place).

  • b. $0.0028 - 0.00012 = 0.0027$ (rounded to four decimal places).

  • c. $984.23 + 10.0 = 994.2$ (rounded to tenths place).

• Multiplication and Division Rule: The result must have the same number of significant figures as the measurement with the fewest significant figures.

  • a. $4.001 \times 15.0 = 60.0$ (3 significant figures).

  • b. $10.50 / 2 = 5$ (1 significant figure based on the number 2, assuming it is a measurement).

  • c. $0.0034 \times 0.0002 = 0.0000007$ (1 significant figure).

Chemical Reactions and the Periodic Table

• Physical vs. Chemical Change:

  • Physical Change: Affects the form of a substance but not its chemical composition. Example: Melting ice.

  • Chemical Change: A process where one or more substances are altered into one or more new and different substances. Example: Rusting iron.

• Evidence of Chemical Reactions:

  • Observations: Gas production (bubbles), color change, formation of a precipitate (solid), or temperature change (release/absorption of energy).

  • Lab Example: Mixing vinegar and baking soda to produce $CO_2$ gas bubbles.

• Periodic Table Organization:

  • Period: A horizontal row. Example: Lithium ($Li$), Beryllium ($Be$), and Boron ($B$).

  • Total Periods: 7.

  • Group: A vertical column. Elements in a group have similar chemical properties. Example: Lithium ($Li$), Sodium ($Na$), and Potassium ($K$).

  • Total Groups: 8

  • Reactivity: Elements in the same group react similarly because they have the same number of valence electrons.

• Valence Electrons and Group Names:

Atomic Theory and Structure

• History of the Atomic Model:

  • Dalton: Proposed the solid sphere model; atoms are indivisible particles.

  • Thomson: Discovered the electron using the cathode ray tube; proposed the Plum Pudding Model.

  • Rutherford: Conducted the Gold Foil Experiment; discovered the dense, positive nucleus and that atoms are mostly empty space.

  • Bohr: Proposed that electrons travel in fixed circular orbits (energy levels) around the nucleus.

• Subatomic Particles:

• Atomic Mathematics:

  • Atomic Number: Always equals the number of protons.

  • Mass Number: Calculated as $\text{Protons} + \text{Neutrons}$.

  • Neutral Atoms: Electrons equal protons if the net charge is zero.

• Ions:

  • Cation: A positively charged ion made by losing electrons. Typically formed by metals (Groups 1 and 2).

  • Anion: A negatively charged ion made by gaining electrons. Typically formed by non-metals (Groups 15, 16, and 17).

• Bohr-Rutherford Examples:

  • Sodium ($Na$): 11 protons, 11 electrons ($2, 8, 1$). Sodium Ion ($Na^+$): 11 protons, 10 electrons ($2, 8$). $Na^+$ is more stable because it has a full valence shell.

  • Oxygen ($O$): 8 protons, 8 electrons ($2, 6$). Oxide Ion ($O^{2-}$): 8 protons, 10 electrons ($2, 8$). $O^{2-}$ is more stable due to the full octet.

• The Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a full outer shell of 8 valence electrons (stable configuration). The "Rule of 0" refers to the net charge of a stable compound being zero.

Nomenclature and Formulas

• Naming Compounds:

  • $Li_2O$: Lithium oxide

  • $K(NO_3)$: Potassium nitrate

  • $HCl$: Hydrochloric acid

  • $Ca_2C$: Calcium carbide

  • $H_3PO_4$: Phosphoric acid

  • $AlF_3$: Aluminum fluoride

  • $(NH_4)Cl$: Ammonium chloride

  • $PbO$: Lead (II) oxide

• Formula Writing:

  • Iron (III) chloride: $FeCl_3$

  • Aluminum phosphate: $AlPO_4$

  • Hydrobromic acid: $HBr$

  • Sodium sulfate: $Na_2SO_4$

  • Copper (II) nitrate: $Cu(NO_3)_2$

  • Calcium oxide: $CaO$

  • Sulfurous acid: $H_2SO_3$

  • Ammonium oxide: $(NH_4)_2O$

Nuclear Chemistry

• Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons.

  • Hyphen Notation: Carbon-14.

  • Nuclear Symbol: $_{6}^{14}C$. (14 is the mass number, 6 is the atomic number).

• Problem: 7 protons, 8 neutrons, 8 electrons:

  • Element is Nitrogen ($N$). Charge is $-1$.

  • Hyphen Notation: Nitrogen-15.

  • Nuclear Symbol: $_{7}^{15}N^-$.

• Fusion vs. Fission:

  • Nuclear Fusion: Small nuclei combine to form a heavier nucleus (occurs in stars).

  • Nuclear Fission: A heavy nucleus splits into smaller nuclei (occurs in nuclear power plants).

• Radioactive Decay:

  • Alpha Decay: Loss of an alpha particle ($_2^4He$). The mass number decreases by 4 and the atomic number decreases by 2.

    • Example: $_{92}^{238}U \rightarrow _{90}^{234}Th + _2^4He$

  • Beta Decay: A neutron turns into a proton and releases an electron ($_{-1}^0e$). The atomic number increases by 1.

    • Example: $_6^{14}C \rightarrow _7^{14}N + _{-1}^0e$

• Stability: Determined by the ratio of neutrons to protons in the nucleus.

Bonding and Molecular Geometry

• Electron Orbitals:

  • s-block: Groups 1, 2, and Helium (2 electrons max).

  • p-block: Groups 13-18 (6 electrons max).

  • d-block: Transition metals (10 electrons max).

  • f-block: Lanthanides and Actinides (14 electrons max).

• Copper Configuration ($Cu$):

  • Full: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^{10}$.

  • Noble Gas: $[Ar]4s^1 3d^{10}$.

• Types of Bonds:

  • Ionic: Electrostatic attraction following the transfer of electrons from a metal to a non-metal. Formulas are simplified ratios (empirical). Example: $NaCl$.

  • Covalent: Sharing of electron pairs between non-metals.

    • Polar Covalent: Unequal sharing.

    • Non-polar Covalent: Equal sharing.

  • Metallic: Attraction between metal cations and a "sea" of delocalized valence electrons.

  • Network Covalent: Atoms bonded in a continuous web (e.g., Diamond).

• Structural Concepts:

  • Coefficient vs. Subscript: A coefficient indicates the number of molecules ($2H_2O$); a subscript indicates the number of atoms within a molecule ($H_2O$).

  • Isomers: Molecules with the same molecular formula but different structural arrangements.

  • HONC 1234 Rule: Hydrogen forms 1 bond, Oxygen forms 2, Nitrogen forms 3, and Carbon forms 4.

• Bonding Electrons: Single bond = 2 electrons; Double bond = 4 electrons.

• Lewis Diagrams and Geometry:

  • $CO_2$: Linear

  • $PCl_3$: Trigonal Pyramidal

  • $SCl_2$: Bent

  • $NH_3$: Trigonal Pyramidal

  • $CCl_4$: Tetrahedral

  • $HCN$: Linear

• Polarity and Dissolution:

  • Water is polar. "Like dissolves like" means water dissolves polar and ionic substances, but not non-polar ones.

  • $CH_4$ (Methane): Lewis diagram shows symmetric distribution; it is non-polar (dipoles cancel).

  • $CH_3F$: Asymmetric; because Fluorine is highly electronegative, it is polar.

• Electronegativity Scale:

  • $0.0 - 0.4$: Non-polar covalent.

  • $0.4 - 1.7$: Polar covalent.

  • > 1.7: Ionic.

Stoichiometry and Yield

• Mole Conversions:

  • 4.5 g of Al: $4.5\,g / 26.98\,g/mol = 0.167\,moles$.

  • 58.3 molecules of CO: $58.3 / (6.02 \times 10^{23}) = 9.68 \times 10^{-23}\,moles$.

  • 5.6 g of $H_2$ gas: $5.6\,g / 2.02\,g/mol = 2.77\,moles \times 22.4\,L/mol = 62.05\,L$ .

• Limiting Reactant Problem ($1SO_2 + 1H_2O \rightarrow 1H_2SO_3$):

  • $12.4\,g\,SO_2 = 0.193\,mol$.

  • $17.9\,g\,H_2O = 0.994\,mol$.

  • Limiting: $SO_2$. Max product: $0.193\,mol \times 82.07\,g/mol = 15.84\,g\,H_2SO_3$.

• Percent Yield ($1Cu + 2AgNO_3 \rightarrow 2Ag + 1Cu(NO_3)_2$):

  • $3.5\,g\,Cu = 0.055\,mol$. Theoretical yield of $Ag = 0.055 \times 2 = 0.11\,mol \times 107.87\,g/mol = 11.87\,g$.

  • Percent Yield: $(10.2 / 11.87) \times 100 = 85.9\%$.

Acids and Bases

• The pH Scale:

  • $0 - 6.9$: Acidic (0-3 is strong, 4-6 is weak).

  • $7.0$: Neutral.

  • $7.1 - 14$: Basic (8-10 is weak, 11-14 is strong).

• Acid/Base Definitions:

  • Arrhenius Base: Produces $OH^-$ in water.

  • Bronsted-Lowery Acid: Proton ($H^+$) donor.

  • Lewis Base: Electron pair donor.

• Naming Rules:

  • Binary Acids: "Hydro-" + base name + "-ic acid" (e.g., $HCl$ = Hydrochloric acid).

  • Oxyacids: "-ate" becomes "-ic"; "-ite" becomes "-ous" (e.g., $H_2SO_4$ from sulfate = Sulfuric acid).

• Calculations:

  • Titration ($HCl$ vs $KOH$): $M_aV_a = M_bV_b \Rightarrow (0.67)(43.6) = M_b(25.9) \Rightarrow M_b = 1.13\,M$.

  • Titration ($HCl$ vs $Ca(OH)_2$): Be aware that $Ca(OH)_2$ has 2 $OH^-$; $M_aV_a = 2M_bV_b \Rightarrow (0.50)(4.6) = 2(M_b)(15.9) \Rightarrow M_b = 0.072\,M$.

  • pH Calculation: $pH = -\log(2.3 \times 10^{-3}) = 2.64$ (Acid).

  • Solution Prep: $M = \frac{n}{V} \Rightarrow 0.05 = \frac{n}{0.4} \Rightarrow 0.02\,mol \times 58.44\,g/mol = 1.17\,g\,NaCl$ .