Advanced Chemistry Second Semester Exam Study Guide 2025 - Comprehensive Notes

First Semester Skills Review

  • Formula Writing: Proficiency in writing formulas for ionic, molecular, acids, hydrocarbons, and alcohols is essential.

  • Balanced Equations: Ability to write complete and balanced chemical equations, including predicting products.

  • Dimensional Analysis & Significant Figures: Correctly perform calculations using dimensional analysis and present answers with the appropriate number of significant figures.

Avogadro’s Number

  • Definition: What is Avogadro's number? 6.022 \times 10^{23}

  • Usage: Know when to apply Avogadro's number in calculations.

  • Example: How many atoms are in 4.80 mol Au?

Periodic Trends

  • Atomic Radius:

    • Trend: Describe the periodic trend for atomic radius and provide the reasoning behind it (increases down and to the left).

    • Ordering: Arrange the following elements in order of increasing atomic radius: Ba, S, Ca, F, Zn.

  • Inverse Relationship: Explain why atomic radius and first ionization energy are inversely related.

  • First Ionization Energy:

    • Definition: Define first ionization energy (the energy required to remove the outermost electron from a neutral atom in the gaseous phase).

    • Trend: Describe the periodic trend for first ionization energy (increases up and to the right).

  • Electronegativity:

    • Definition: Define electronegativity (the ability of an atom to attract bonding electrons in a chemical bond).

    • Trend: Describe the periodic trend for electronegativity (increases up and to the right).

  • Shielding Effect: Explain the shielding effect (the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell).

  • Lone Pair: What is a lone (unshared) pair of electrons?

  • Bond Types: How many electrons are in single, double, and triple bonds? (2, 4, and 6, respectively).

Bonding

  • Bond Type Prediction: What type of bond do phosphorus and bromine form? Why? (covalent, small electronegativity difference).

  • Ionic Bond: Define ionic bond (bond formed through electrostatic attraction between oppositely charged ions).

  • Bonding Electrons: Which electrons are involved in bonding? (valence electrons).

  • Bond Types Comparison: List the three main types of bonds discussed (ionic, covalent, metallic), explain the role of electrons in each type, and indicate which types of elements (metals or nonmetals) they bond together.

  • Metallic Bond: Which of the three types of bonds discussed does not bond atoms to form a new compound, instead forming a “sea of electrons”? (metallic).

Percent Composition

  • Calculation: How is the percent by mass (percent composition) of each element in a compound calculated?
    \text{Percent by mass} = \frac{\text{mass of element}}{\text{mass of compound}} \times 100

  • Application: Be able to use percent composition to determine the mass of a given element in a certain amount of a compound.

Empirical Formulas

  • Reduction: How do you reduce molecular formulas to find empirical formulas? (Divide by the greatest common factor).

  • Determination from % Composition: How do you determine empirical formulas from percent composition? (Convert % to mass, mass to moles, divide by smallest, multiply to whole numbers).

  • Example: What is the empirical formula of a compound that is 32.38% Na, 22.65% S, and 44.99% O? (Na2SO3)

  • Molecular Formula Determination: How do you determine molecular formula given empirical formula and molecular formula mass?
    \text{Molecular formula} = (\text{empirical formula})_n

    where n = \frac{\text{molecular formula mass}}{\text{empirical formula mass}}

  • Example: What is the molecular formula of a compound with the empirical formula PCl5 and molar mass 132.84 g/mol? (It should be noted that the molar mass given in this question is wrong. Using the atomic weights of P (30.97 g/mol) and Cl (35.45 g/mol), the molar mass for PCl5 is actually 208.23 g/mol, and so the molecular formula would be the same as the empirical formula PCl5. If the molar mass was 104.115, the empirical formula would be P{0.5}Cl_{2.5} which is not possible as empirical formulas must only have whole number subscripts.).

Stoichiometry, Limiting Reactants, Excess Reactants, and Percent Yield

  • Definitions: Define stoichiometry and mole ratio.

  • Mole Conversions: Be able to convert from moles of one substance to moles of another using the mole ratio from the balanced chemical equation.

  • Mole-Mass Conversions: Be able to convert from moles of one substance to mass of another using the mole ratio from the balanced chemical equation and the molar mass. Also, you need to know how to convert from mass of one substance to moles of another using the mole ratio from the balanced chemical equation and the molar mass.

  • Mass-Mass Conversions: Be able to convert from mass of one substance to mass of another using the mole ratio from the balanced chemical equation and the molar masses of both substances.

  • Definitions: Define theoretical yield, limiting reactant, and excess reactant.

  • Limiting Reactant & Theoretical Yield: Be able to calculate the theoretical yield and then determine which reactant is the limiting reactant.

  • Excess Reactant Remaining: Be able to determine the amount of excess reactant remaining after the limiting reactant is used up (completely reacted).

  • Definitions: Define actual yield and percent yield.

  • Percent Yield Calculation: Be able to calculate percent yield. \text{Percent Yield} = (\frac{\text{Actual Yield}}{\text{Theoretical Yield}}) \times 100

States of Matter

  • Kinetic Theory: Use the kinetic theory to explain the difference between the particles in a gas and those in a liquid. What are the five assumptions of the kinetic molecular theory of gases?

  • Gas Stoichiometry: What volume of HCl gas is required to react with excess aluminum metal to produce 8.62 L of hydrogen gas at STP?

  • Ideal Gas: What is an ideal gas?

  • Phase Diagram: Find a phase diagram for water. What happens to water at 105°C and 1.00 atm if the pressure increased to 1.5 atm? (water turns from gas to liquid)

  • Compressibility: Which are more compressible – solids, liquids, or gases? (gases).

Gas Laws

Conversion factors and list of various gas laws will be provided.

  • Charles's Law Example: If a sample of gas occupies 7.6 L at 327°C, what will its volume be at 56°C if pressure is constant? Which law does this represent? (Charles's Law).

  • Boyle's Law Example: The pressure on 4.5 L of gas changes from 150 kPa to 75 kPa. What is the new volume in milliliters if temperature is constant? Which law does this represent? (Boyle's Law).

  • Gay-Lussac's Law Example: On a cold winter morning when the temperature is -13°C, the air pressure in an automobile tire is 1.5 atm. If the volume does not change, then what is the pressure after the tire has warmed to 15°C?

  • Combined Gas Law Example: A gas has a pressure of 1.2 atm, a volume of 135 mL, and a temperature of 14°C. After a drop in pressure to 1.0 atm and a drop in volume to 95 mL, what is the new temperature?

  • Ideal Gas Law Equation: What is the equation for the Ideal Gas Law, and what does each variable represent? \textbf{PV = nRT}

  • Ideal Gas Law Example: What is the temperature of a 15.2 mole sample of helium that has a volume of 225 mL and a pressure of 805 torr?

  • Pressure Conversion: Convert 745 mmHg to torr (745 torr).

  • Pressure Conversion: Convert 0.75 atm to mmHg (570 mmHg).

  • Ideal Gas Law with Moles and Mass: A sample of chlorine gas has a volume of 11.5 L, and temperature of 27 °C, and a pressure of 102 kPa. How many moles of gas are there? What is its mass?

  • STP Definition: What is STP? Give values for variables. (Standard Temperature and Pressure: 0 °C and 1 atm)

  • Molar Volume at STP: What volume does 1 mole of any gas occupy at STP? (22.4 L).

  • Gay-Lussac's Law of Combining Volumes: If 0.5 L of oxygen gas reacts with excess hydrogen gas to produce 1.0 L of water vapor, then what volume of water vapor would be produced when 3.5 L of oxygen gas reacts with excess hydrogen gas?

  • Molar Volume at STP Example: What volume will 34g of oxygen occupy at STP?

  • Effusion: Which effuses faster, helium or neon gas? How much faster?

  • Kinetic Energy and Temperature: Explain the relationship between kinetic energy and temperature.

  • Molar Mass Calculation: At SPT 1.92g of a gas has a volume of 100 L. What is the molar mass of the gas?

  • Partial Pressures: A container with pressure 2.34 atm contains oxygen, nitrogen, and hydrogen gases. The partial pressures of oxygen and nitrogen are 0.56 atm and 1.03 atm, respectively. What is the partial pressure of hydrogen? What law do these problems represent? (Dalton's Law of Partial Pressures).

Thermochemistry

  • Heat vs. Temperature: What is the difference between heat and temperature?

  • Standard State: What is meant by standard state?

  • Specific Heat Units: What are the units for specific heat? How do they relate to the definition of specific heat? \frac{J}{g \cdot °C} (energy required to raise 1 gram of a substance 1 degree Celsius).

  • Absolute Zero: What is the temperature at absolute zero? What does this mean? (-273.15 °C, 0 K, the point at which all molecular motion stops).

  • Heat Units: What units are used to measure heat? How many joules equal 56.4 cal?

  • Heat Calculation Example: How many joules are gained when 150 g water goes from 24.4°C to 98.2°C?

  • Specific Heat Calculation Example: A 4.0 g sample of iron was heated from 0.00°C to 20.0°C. It absorbed 35.2 J of heat. What is the specific heat of iron?

  • Energy Absorption Calculation: How much energy does 8.0 g Cu absorb when it is heated from 10.0° to 40.0°C if the specific heat of Cu is 0.384J/(g°C)?

  • Heat of Formation: Define heat of formation (the change in enthalpy during the formation of one mole of the substance from its constituent elements, with all reactants and products in their standard states).

  • Heat of Reaction Calculation: Use heat of formation values from your book to calculate heat of reaction for C2H2(g) + O2(g) \rightarrow CO2(g) + H_2O(g).

  • Exothermic vs. Endothermic: Is the reaction in the problem above exothermic or endothermic? How do you know?

  • Reaction Pathway Graph: Draw the reaction pathway graph for the problem above.

  • Heat Stoichiometry: Be able to do heat stoichiometry problems. What is the change in heat when 4.35g C2H2 combusts?

  • Heats of Phase Change: Define heat of combustion, heat of solidification, heat of vaporization, and heat of fusion, and calculations involving these terms.

  • Phase Changes: Define sublimation, condensation, vaporization, and solidification.

  • Phase Diagrams: Be able to interpret phase diagrams.

  • R-Constant Lab:* -Use a data table (from R-constant lab) to make calculations and draw conclusions.*

Practice Problems

  • Problem 1: A solution of 100.0 mL of 0.200M KOH is mixed with a solution of 200.0 mL of 0.150 M Fe2(SO4)3 and the following reaction takes place: 6KOH(aq) + Fe2(SO4)3 (aq) \rightarrow 3K2SO4(aq) + 2Fe(OH)_3(s). How many grams of the precipitate (solid) form?

  • Problem 2: The metabolic reaction of glucose, C6H{12}O6, in our bodies produces CO2, which is exhaled from our lungs as a gas: C6H{12}O6(aq) + 6O2(g) \rightarrow 6CO2(g) + 6H2O(l). Our bodies produce about 24.0 L of oxygen at normal body conditions (37°C, 0.970 atm.) How many liters of oxygen will react at STP when 24.5 g of glucose is consumed? How many oxygen atoms are used in the reaction?

  • Problem 3: When solutions containing lead (II) ions and chloride ions are mixed, lead (II) chloride precipitates: Pb^{2+} (aq) + 2Cl^-(aq) \rightarrow PbCl2(s) \;\; \Delta H = -359.4 \text{ kJ}. Calculate \Delta H for the precipitation of PbCl2(s) if 5.00 g Pb^{2+} ions and 5.00 g Cl^- ions are present.

  • Problem 4: When 5.0g hydrogen and 28.0g iron (II) oxide react, how many grams of solid iron form?

  • Problem 5: How many molecules are in a 9.00L balloon at STP filled with carbon dioxide? Would your answer change if the gas was carbon monoxide?

  • Problem 6: How many liters of ammonia form when 6.0L nitrogen and 10.0L hydrogen react?

  • Problem 7: Acetylene gas, C2H2, is produced by adding water to calcium carbide, CaC2: CaC2 + 2H2O \rightarrow C2H2 + Ca(OH)2. What volume of acetylene is produced when 50.80g calcium carbide react at 65°C and 654 torr?

  • Problem 8: A 3.0L flexible container of gas at 24°C was heated to 96°C. What is the new volume? What gas law describes this? (Charles's Law)

  • Problem 9: What is the empirical formula of a compound that is 25.9% nitrogen and 74.1% oxygen?

  • Problem 10: How many joules are gained when 150 g water goes from 24.4°C to 98.2°C?

  • Problem 11: A 4.0 g sample of iron was heated from 0.00°C to 20.0°C. It absorbed 35.2 J of heat. What is the specific heat of iron?

  • Problem 12: How much energy does 8.0 g Cu absorb when it is heated from 10.0° to 40.0°C if the specific heat of Cu is 0.384J/(g°C)?

  • Problem 13: A container with pressure 2.34 atm contains oxygen, nitrogen, and hydrogen gases. The partial pressures of oxygen and nitrogen are 0.56 atm and 1.03 atm, respectively. What is the partial pressure of hydrogen? What law do this problems represent? (Dalton's Law of Partial Pressures)

  • Problem 14: What volume of HCl gas is required to react with excess aluminum metal to produce 8.62 L of hydrogen gas at STP?

  • Problem 15: What are the five assumptions of the kinetic molecular theory of gases?

  • Problem 16: What is an ideal gas?

  • Problem 17: Indicate which element in each of the following pairs has the greater first ionization energy and greater atomic radius.

    • a) Li or B

    • b) Mg or Sr

    • c) Cs or Al

  • Problem 18: Which has the greater third ionization energy; Mg or Al? Why?

  • Problem 19: Which has a greater ionic radius; Mg^{2+} or Al^{3+}?

  • Problem 20: Use heats of formation to determine the heat change for the following reaction. $$CaSO4 + BaCl2 \rightarrow Ba