Honors Chemistry Semester 2 Final Exam Study Guide Flashcards

Chemical Equations and Symbols

• State Symbols in Chemical Equations: These symbols indicate the physical state of each substance in a reaction:

  • $(s)$: Denotes a solid substance.

  • $(l)$: Denotes a liquid substance.

  • $(g)$: Denotes a gaseous substance.

  • $(aq)$: Denotes an aqueous solution, meaning the substance is dissolved in water.

• Reactants vs. Products:

  • Reactants: The starting materials in a chemical reaction, typically written on the left side of the equation.

  • Products: The new substances formed as a result of the reaction, written on the right side of the equation.

• The Yield Arrow: The arrow symbol ($\rightarrow$) represents the chemical change occurring. It is read as "yields," "produces," or "reacts to form."

• Law of Conservation of Mass: Chemical equations must be balanced to satisfy this law, which states that matter cannot be created or destroyed. The number of atoms for each element must be identical on both the reactant and product sides.

• Balancing Mechanics:

  • Coefficients: These are the large numbers placed in front of chemical formulas (e.g., $2H_2O$). You can and must change these to balance the equation.

  • Subscripts: These are the small numbers within a formula indicating the number of atoms in a molecule (e.g., the "2" in $H_2O$). You must never change subscripts, as doing so changes the identity of the substance.

Balancing Practice and Reaction Types

• Balancing Examples:

  • Ethane Combustion: $2C_2H_6 + 7O_2 \rightarrow 4CO_2 + 6H_2O$

  • Iron Oxidation: $4Fe + 3O_2 \rightarrow 2Fe_2O_3$

  • Double Replacement: $AlCl_3 + 3NaOH \rightarrow Al(OH)_3 + 3NaCl$

  • Single Replacement: $Mg + 2HCl \rightarrow MgCl_2 + H_2$

• The Five Major Reaction Types:

  1. Synthesis: Two or more substances combine to form a single product ($A + B \rightarrow AB$).

  2. Decomposition: A single reactant breaks down into two or more simpler products ($AB \rightarrow A + B$).

  3. Single Replacement: One element replaces another element in a compound ($A + BC \rightarrow AC + B$). A metal replaces a metal, or a nonmetal replaces a nonmetal.

  4. Double Replacement: The ions of two compounds exchange places in an aqueous solution to form two new compounds ($AB + CD \rightarrow AD + CB$).

  5. Combustion: A hydrocarbon reacts with oxygen gas ($O_2$) to produce carbon dioxide ($CO_2$) and water vapor ($H_2O$). This reaction is characterized by the release of significant energy as heat and light.

Solubility Rules and the Activity Series

• Predicting Precipitates with Solubility Charts: A solubility chart identifies whether an ionic compound is soluble (dissolves) or insoluble (forms a solid precipitate) in water. If a product of a double replacement reaction is marked as insoluble ($s$), a precipitate forms.

• Solubility Application:

  • Silver Chloride ($AgCl$): According to solubility rules, most chloride ($Cl^-$) salts are soluble, but silver ($Ag^+$) is a primary exception. Therefore, $AgCl$ is insoluble.

  • Barium Sulfate ($BaSO_4$): Most sulfates are soluble, but Barium ($Ba^{2+}$) is an exception, making barium sulfate insoluble.

• Activity Series in Single Replacement: This list ranks metals by their reactivity. A metal can only replace another metal in a compound if it is higher (more reactive) on the activity series.

  • Example 1: Copper ($Cu$) will not replace Zinc ($Zn$) in zinc chloride because copper is lower on the activity series than zinc.

  • Example 2: Magnesium ($Mg$) will replace Copper ($Cu$) in copper(II) sulfate because magnesium is higher on the activity series than copper.

Stoichiometry and Molar Mass

• Molar Mass: This is the mass of one mole of a substance ($g\,mol^{-1}$). It is calculated by summing the atomic masses of all atoms in the chemical formula, found on the Periodic Table.

  • Molar Mass of $CO_2$: $12.01 + 2 \times 16.00 = 44.01\,g\,mol^{-1}$.

  • Molar Mass of $Ca(OH)_2$: $40.08 + 2 \times (16.00 + 1.01) = 74.10\,g\,mol^{-1}$.

• Molar Conversions:

  • Grams to Moles: Divide mass by molar mass ($18.0\,g\,H_2O / 18.02\,g\,mol^{-1} = 1.00\,mol$).

  • Moles to Grams: Multiply moles by molar mass ($3.50\,mol\,CO_2 \times 44.01\,g\,mol^{-1} = 154\,g$).

• Mole Ratios: These are conversion factors derived from the coefficients of a balanced chemical equation. They show the relative amounts of reactants and products.

  • Example: In $2Al + 3Cl_2 \rightarrow 2AlCl_3$, the mole ratio between $Cl_2$ and $AlCl_3$ is $\frac{3\,mol\,Cl_2}{2\,mol\,AlCl_3}$.

• Standard Temperature and Pressure (STP):

  • Definition: At STP ($0^\circ\text{C}$, $1\,atm$), $1\,mol$ of any gas occupies exactly $22.4\,dm^3$.

Limiting Reactants and Percent Yield

• Limiting vs. Excess Reactants:

  • Limiting Reactant: The reactant that is completely consumed first. it limits the amount of product that can form.

  • Excess Reactant: The reactant that remains after the reaction stops.

• Determining the Limiting Reactant:

  1. Convert both starting masses to moles of the same product.

  2. The reactant that produces the smaller amount of product is the limiting reactant.

• Theoretical Yield: The maximum amount of product that can be produced, as calculated through stoichiometry using the limiting reactant.

• Percent Yield: A measure of the efficiency of a reaction.

  • Formula: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100$

  • Calculation Example: If actual is $24.8\,g$ and theoretical is $32.0\,g$, the yield is $\frac{24.8}{32.0} \times 100 = 77.5\%$.

Thermochemistry and Energy Changes

• Exothermic Reactions: Reactions that release heat to the surroundings. The enthalpy change ($\Delta H$) is negative.

• Endothermic Reactions: Reactions that absorb heat from the surroundings. The enthalpy change ($\Delta H$) is positive.

• Energy Conversions: $1\,kJ = 1000\,J$.

• The Heat Equation: $q = mc\Delta T$

  • $q$: Heat energy (Joules).

  • $m$: Mass (grams).

  • $c$: Specific heat capacity ($J\,g^{-1}\,^\circ\text{C}^{-1}$).

  • $\Delta T$: Change in temperature ($T_{\text{final}} - T_{\text{initial}}$).

• Heating Curves:

  • Slanted Sections: Represent temperature changes where kinetic energy is increasing.

  • Flat Sections (Plateaus): Represent phase changes (melting or boiling). Temperature remains constant because the energy added is being used to overcome intermolecular forces rather than increasing particle speed.

Hess's Law and Bond Enthalpy

• Hess's Law: The total enthalpy change for a reaction is the same regardless of the number of steps taken.

  • Rule 1: If you reverse an equation, the sign of $\Delta H$ is flipped ($+$ becomes $-$).

  • Rule 2: If you multiply the coefficients by a factor, you must multiply $\Delta H$ by that same factor.

• Bond Enthalpy: The energy required to break a chemical bond.

  • Bonds Broken: Energy is absorbed (endothermic process).

  • Bonds Formed: Energy is released (exothermic process).

  • Calculation Formula: $\Delta H = \sum (\text{Bonds Broken}) - \sum (\text{Bonds Formed})$.

Atomic Structure and Lewis Structures

• Valence Electrons: The electrons in the outermost energy level of an atom, responsible for chemical bonding.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a full outer shell of 8 electrons (mimicking Noble Gases).

• Pairs of Electrons:

  • Shared Pair: A pair of electrons involved in a covalent bond between atoms.

  • Lone Pair: A pair of valence electrons not involved in bonding.

• Bond Types and Strengths:

  • Single Bond: 1 shared pair ($2\,e^-$).

  • Double Bond: 2 shared pairs ($4\,e^-$).

  • Triple Bond: 3 shared pairs ($6\,e^-$).

Kinetic Molecular Theory (KMT) and Gas Laws

• KMT Principles: Gas particles are in constant, rapid, random motion. Their collisions are elastic, and they have negligible volume compared to the space between them.

• Temperature and Motion: Average kinetic energy is directly proportional to the Kelvin temperature. If the Kelvin temperature doubles, the average kinetic energy doubles.

• Absolute Temperature: Gas law calculations must use Kelvin to avoid zero or negative denominators/results.

  • Conversion: $K = ^\circ\text{C} + 273.15$.

• Gas Relationships:

  • Boyle’s Law ($P_1V_1 = P_2V_2$): Pressure and volume are inversely proportional (as volume decreases, pressure increases).

  • Charles’s Law ($\frac{V_1}{T_1} = \frac{V_2}{T_2}$): Volume and Kelvin temperature are directly proportional.

  • Gay-Lussac’s Law ($\frac{P_1}{T_1} = \frac{P_2}{T_2}$): Pressure and Kelvin temperature are directly proportional.

  • Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$.

• Ideal Gas Law: $PV = nRT$

  • $P$: Pressure ($atm$, $kPa$).

  • $V$: Volume ($dm^3$).

  • $n$: Moles.

  • $R$: Ideal gas constant (e.g., $0.0821\,dm^3\,atm\,mol^{-1}\,K^{-1}$).

  • $T$: Kelvin Temperature.

Dalton's and Graham's Laws

• Dalton’s Law of Partial Pressures: The total pressure of a gas mixture is the sum of the partial pressures of the individual gases.

  • $P_{\text{total}} = P_{\text{gas}} + P_{\text{water vapor}}$.

  • When collecting gas over water, the "dry gas" pressure is found by subtracting the water vapor pressure from the total barometric pressure.

• Graham’s Law of Effusion: Compares the rate at which gases escape through small holes. Lighter gases (lower molar mass) effuse faster than heavier gases.

  • $\text{Rate Ratio} = \sqrt{\frac{\text{Molar Mass}_B}{\text{Molar Mass}_A}}$.

• Phase Diagrams:

  • Triple Point: The specific temperature and pressure where all three phases (solid, liquid, gas) coexist in equilibrium.

  • Boiling Point: Water can boil below $100^\circ\text{C}$ if the surrounding atmospheric pressure is lowered (e.g., at high altitudes).

Solutions and Molarity

• Solution Components:

  • Solute: The substance being dissolved (e.g., salt).

  • Solvent: The substance doing the dissolving (e.g., water).

• Saturation Levels:

  • Saturated: Contains the maximum amount of dissolved solute for a given temperature.

  • Unsaturated: Contains less than the maximum amount of solute.

  • Supersaturated: Contains more than the theoretical maximum; highly unstable.

• Gas Solubility Factors:

  • Increasing temperature decreases gas solubility in liquids.

  • Decreasing pressure decreases gas solubility (why soda goes flat when the cap is removed).

• Molarity Formula: $M = \frac{\text{moles of solute}}{\text{dm}^3\text{ of solution}}$.

Acids, Bases, and pH

• Definitions:

  • Acid: A substance that produces $H^+$ ions in solution (pH < 7).

  • Base: A substance that produces $OH^-$ ions in solution (pH > 7).

  • Strong vs. Weak: Strong acids dissociate 100% in water; weak acids only partially dissociate.

• Naming Acids:

  • $HCl$: Hydrochloric acid.

  • $HNO_3$: Nitric acid.

  • $H_2SO_4$: Sulfuric acid.

• pH and pOH Calculations:

  • $pH = -\log([H^+])$

  • $pOH = -\log([OH^-])$

  • $pH + pOH = 14.00$ (at $25^\circ\text{C}$).

  • $[H^+][OH^-] = 1.0 \times 10^{-14}$.

• Neutralization Reactions: A reaction between an acid and a base that produces water ($H_2O$) and an ionic salt.

Extended Response Procedure: Combustion Analysis

For any combustion lab (Ethane, Propane, Methanol), follow these exhaustive steps:

• A. Balanced Equation: Ensure the correct ratio of fuel to $O_2$ yields $CO_2$ and $H_2O$.

• B. Reaction Type: Always Combustion.

• C. Lewis Structures:

  • $O_2$ has a double bond.

  • $CO_2$ has two double bonds.

  • $H_2O$ has two single bonds and two lone pairs on Oxygen.

• D. Bond Counting: Multiply the bonds in one molecule by the coefficient in the balanced equation.

• E. Enthalpy ($\Delta H$): sum the energies of all reactant bonds (broken) and subtract the sum of all product bonds (formed).

• F. Limiting Reactant: Convert initial grams of fuel and grams of oxygen to grams of $CO_2$. The lower result identifies the limiting reactant.

• G. Theoretical Yield: The mass calculated from the limiting reactant in step F.

• H. Percent Yield: Divide the "Actual mass collected" provided in the prompt by the result from step G.

• I. Dalton's Law: Subtract the water vapor pressure from the total pressure to find the partial pressure of the dry $CO_2$.

• J. Ideal Gas Law: Use $V = \frac{nRT}{P}$ where $n$ is the theoretical moles of $CO_2$, $T$ is in Kelvin, and $P$ is the dry gas pressure from step I.

• K. pH Calculation: Use $pH = -\log([H^+])$. If pH is < 7, the solution is acidic.

• L. Effusion Analysis: Use Graham's Law to compare the molar mass of $CO_2$ ($44.01\,g\,mol^{-1}$) to $O_2$ ($32.00\,g\,mol^{-1}$).