Comprehensive Study Guide for CBA and Chemistry Final Exam

Stoichiometry and Chemical Equations (30%)

  • Balancing Chemical Equations

    • Chemical equations must be balanced to satisfy the Law of Conservation of Mass, which states that matter cannot be created or destroyed in a chemical reaction.

    • The number of atoms for each element on the reactant side must equal the number of atoms for each element on the product side.

    • Subscripts define the identity of the substance and cannot be changed. Only coefficients (the numbers in front of chemical formulas) are adjusted to balance the equation.

  • Molar Ratios

    • Molar ratios are conversion factors derived from the coefficients of a balanced chemical equation.

    • These ratios are used to relate the number of moles of one substance to the number of moles of another substance involved in the reaction.

    • Example: For the reaction 2H2+O2→2H2O2H_2 + O_2 \rightarrow 2H_2O, the molar ratio between H2H_2 and O2O_2 is 2 moles H21 mole O2\frac{2\text{ moles } H_2}{1\text{ mole } O_2}.

  • Molar Mass

    • Molar mass is the mass of one mole of a substance, typically expressed in units of g mol−1g\,mol^{-1}.

    • It is calculated by summing the atomic masses of all atoms present in the chemical formula as found on the periodic table.

    • Example: The molar mass of H2OH_2O is calculated as (2×1.01)+(1×16.00)=18.02 g mol−1(2 \times 1.01) + (1 \times 16.00) = 18.02\,g\,mol^{-1}.

  • Percent Composition

    • Percent composition is the percentage by mass of each element in a compound.

    • The formula used is:     % element=mass of element in 1 mole of compoundmolar mass of compound×100\% \text{ element} = \frac{\text{mass of element in 1 mole of compound}}{\text{molar mass of compound}} \times 100

  • Atoms from Mole Calculations

    • Avogadro's Number is the constant used to convert between moles and particles (atoms, molecules, or formula units): 6.02×10236.02 \times 10^{23}.

    • To find the number of atoms from a given mass:

      1. Convert grams to moles using molar mass: moles=massmolar mass\text{moles} = \frac{\text{mass}}{\text{molar mass}}.

  1. Convert moles to atoms using Avogadro's number: atoms=moles×(6.02×1023)\text{atoms} = \text{moles} \times (6.02 \times 10^{23}) .

Empirical and Molecular Formulas (5%)

  • Empirical Formula

    • The empirical formula represents the simplest whole-number ratio of the atoms of each element present in a compound.

    • Example: The empirical formula for glucose (C6H12O6C_6H_{12}O_6) is CH2OCH_2O.

  • Molecular Formula

    • The molecular formula represents the actual number of atoms of each element in one molecule of the substance.

    • The molecular formula is always a whole-number multiple of the empirical formula:     Molecular Formula=n×(Empirical Formula)\text{Molecular Formula} = n \times (\text{Empirical Formula})

    • The multiplier nn is found by: n=Molar Mass of Molecular FormulaMolar Mass of Empirical Formulan = \frac{\text{Molar Mass of Molecular Formula}}{\text{Molar Mass of Empirical Formula}}.

Gas Laws and Solution Dilution (10%)

  • Gas Law Calculations

    • Ideal Gas Law: PV=nRTPV = nRT

      • PP = Pressure

      • VV = Volume (typically in dm3dm^3)

      • nn = Number of moles

      • RR = Ideal gas constant

      • TT = Temperature (must be in Kelvin: K=∘C+273K = ^\circ C + 273)

    • Combined Gas Law: P1V1T1=P2V2T2\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}

  • Solution Dilution

    • Dilution is the process of reducing the concentration of a solute in a solution, usually by adding more solvent.

    • The dilution equation is: M1V1=M2V2M_1V_1 = M_2V_2

      • M1M_1 = Initial Molarity (mol dm−3mol\,dm^{-3})

      • V1V_1 = Initial Volume (dm3dm^3 or cm3cm^3)

      • M2M_2 = Final Molarity (mol dm−3mol\,dm^{-3})

      • V2V_2 = Final Volume (dm3dm^3 or cm3cm^3)

Solubility Trends (5%)

  • Solute/Solvent Combinations

    • Solubility follows the general principle of "Like Dissolves Like":

      • Polar solutes dissolve in polar solvents (e.g., salt in water).

      • Non-polar solutes dissolve in non-polar solvents (e.g., oil in benzene).

  • Temperature and Pressure Effects

    • Solids in Liquids: Generally, solubility increases as temperature increases.

    • Gases in Liquids: Solubility decreases as temperature increases and increases as the partial pressure of the gas above the liquid increases (Henry's Law).

Acid and Base Properties (30%)

  • Arrhenius Acid and Base Definitions

    • Arrhenius Acid: A substance that increases the concentration of hydrogen ions (H+H^+) when dissolved in water.

    • Arrhenius Base: A substance that increases the concentration of hydroxide ions (OH−OH^-) when dissolved in water.

  • Bronsted-Lowry Acid and Base Definitions

    • Bronsted-Lowry Acid: A proton (H+H^+) donor.

    • Bronsted-Lowry Base: A proton (H+H^+) acceptor.

    • This theory introduces Conjugate Acid-Base Pairs: When an acid donates a proton, it becomes a conjugate base; when a base accepts a proton, it becomes a conjugate acid.

  • pH Calculations and Definitions

    • pH is a measure of the acidity or basicity of an aqueous solution.

    • The pH scale typically ranges from 0 to 14, where 7 is neutral, less than 7 is acidic, and greater than 7 is basic.

    • Formula for pH: pH=−log⁡[H+]pH = -\log[H^+]

    • Formula for pOH: pOH=−log⁡[OH−]pOH = -\log[OH^-]

    • Relationship between pH and pOH: pH+pOH=14pH + pOH = 14

    • Calculating ion concentration from pH: [H+]=10−pH[H^+] = 10^{-pH}

  • Acid-Base Reactants and Products

    • Neutralization Reactions: When a strong acid reacts with a strong base, the typical products are a salt and water.

    • General Equation: HA+BOH→BA+H2OHA + BOH \rightarrow BA + H_2O

Thermochemistry Concepts (15%)

  • Energy Definitions

    • Kinetic Energy: The energy an object possesses due to its motion.

    • Potential Energy: Stored energy based on the position, arrangement, or state of an object (e.g., chemical potential energy stored in bonds).

  • Specific Heat Calculations

    • Specific heat capacity (cc) is the amount of heat required to raise the temperature of one gram of a substance by one degree Celsius.

    • The heat transfer formula is: q=mcΔTq = mc\Delta T

      • qq = Heat energy (Joules, JJ)

      • mm = Mass (grams, gg)

      • cc = Specific heat capacity (J g−1 ∘C−1J\,g^{-1}\,^\circ C^{-1})

      • ΔT\Delta T = Change in temperature (Tfinal−TinitialT_{final} - T_{initial}, in  ∘C\text{ }^\circ C)

  • Exothermic and Endothermic Definitions

    • Exothermic: A process that releases heat to its surroundings. The enthalpy change (ΔH\Delta H) is negative, and the surroundings feel warmer.

    • Endothermic: A process that absorbs heat from its surroundings. The enthalpy change (ΔH\Delta H) is positive, and the surroundings feel cooler.

Nuclear Chemistry (5%)

  • Nuclear Radiation Types

    • Alpha Particle (α\alpha): Helium nucleus (24He^4_2He), low penetrating power, high ionizing power.

    • Beta Particle (β\beta): High-speed electron (−10e^0_{-1}e), moderate penetrating power.

    • Gamma Ray (γ\gamma): High-energy electromagnetic radiation, very high penetrating power.

  • Radioactive Decay

    • The process by which an unstable atomic nucleus loses energy by emitting radiation.

    • Decay processes result in the transformation of an atom into a different element or isotope.

  • Fusion and Fission

    • Nuclear Fission: The splitting of a heavy, unstable nucleus into two or more lighter nuclei, releasing a large amount of energy (utilized in nuclear power plants).

    • Nuclear Fusion: The process where two light nuclei combine to form a heavier nucleus, releasing massive amounts of energy (the process that powers stars/the Sun).