AP Chemistry Midterm Study Guide: Atomic Structure, Bonding, Gases, Thermochemistry, and Chemical Reactions

Unit 1: Atomic Structure and Properties

1.1 Moles and Molar Mass

The mole is a fundamental unit in chemistry for counting particles, and molar mass relates mass to the number of moles.

• Mole: The amount of substance containing as many entities as there are atoms in 12 g of carbon-12 (Avogadro's number: $6.022 \times 10^{23}$).

• Molar Mass: The mass of one mole of a substance, expressed in g/mol.

• Example: Calculate the number of oxygen atoms in 3.00 g of CO2.

1.2 Mass Spectroscopy of Elements

Mass spectroscopy is used to determine isotopic masses and abundances, allowing calculation of average atomic mass.

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

• Average Atomic Mass: Weighted average of isotopic masses.

• Example: Given isotopic abundances, calculate the average atomic mass.

1.3 Elemental Composition of Pure Substances

Determining the percent composition of elements in compounds is essential for chemical analysis.

• Percent Composition: $(\text{Percent by mass}) = \frac{\text{Mass of element in 1 mol of compound}}{\text{Molar mass of compound}} \times 100\%$

• Example: Calculate the percent by mass of nitrogen in NaNO3.

1.4 Composition of Mixtures

Mixtures contain more than one substance, and their composition can be described by mass percent.

• Mass Percent: $(\text{Mass percent}) = \frac{\text{Mass of component}}{\text{Total mass of mixture}} \times 100\%$

• Example: Calculate the mass percent of NaCl in a solution containing 12 g NaCl dissolved in 88 g water.

1.5 Atomic Structure and Electron Configuration

Electron configuration describes the arrangement of electrons in atoms.

• Ground-State Electron Configuration: The lowest energy arrangement of electrons in an atom.

• Example: Write the ground-state electron configuration for a sodium atom.

1.6 Photoelectron Spectroscopy (PES)

PES is used to determine the energies of electrons in atoms and ions.

• PES Spectrum: Shows peaks corresponding to electrons in different energy levels.

• Example: Explain why core electrons have higher binding energies than valence electrons.

1.7 Periodic Trends

Periodic trends describe how properties of elements change across the periodic table.

• Atomic Radius: Decreases across a period, increases down a group.

• Ionization Energy: Increases across a period, decreases down a group.

• Electronegativity: Increases across a period, decreases down a group.

• Example: Which element has the largest atomic radius: Na, Mg, Al, or Cl?

1.8 Electromagnetic Radiation, Photons, and Electron Transitions

Atoms absorb or emit energy as photons during electron transitions.

• Energy of a Photon: $E = h\nu = \frac{hc}{\lambda}$

• Example: Calculate the wavelength of a photon released when an electron transitions from $n=2$ to $n=1$ and releases $5.42 \times 10^{-19}$ J.

Unit 2: Molecular and Ionic Compound Structure and Properties

2.1 Types of Chemical Bonds

Chemical bonds include ionic, covalent, and metallic bonds, each with distinct properties.

• Ionic Bonds: Formed by transfer of electrons between metals and nonmetals.

• Covalent Bonds: Formed by sharing electrons between nonmetals.

• Example: Classify bonds as ionic or covalent and predict properties.

2.2 Structure of Ionic Solids

Ionic solids are held together by electrostatic forces and have high melting points.

• Lattice Energy: Energy required to separate one mole of an ionic solid into gaseous ions.

• Example: Which compound has the greatest lattice energy: NaCl, MgO, or CaBr2?

2.4 Structure of Covalent Molecules

Covalent molecules can be represented using Lewis structures to show bonding and lone pairs.

• Lewis Structure: Diagram showing the arrangement of electrons in a molecule.

• Example: Draw the Lewis structure for SO2 and determine the bond order on sulfate.

2.5 Lewis Diagrams

Lewis diagrams help evaluate resonance and formal charge to determine the best structure.

• Resonance: Occurs when more than one valid Lewis structure can be drawn for a molecule.

• Example: Two Lewis structures are proposed for NO2; which is more stable and why?

Unit 3A: Gases (Gas Laws & Kinetic Molecular Theory)

3A.1 Pressure and Temperature

Gas pressure and temperature are related to the kinetic energy of particles.

• Average Kinetic Energy: $KE_{avg} = \frac{3}{2}RT$

• Example: Describe what happens to the average kinetic energy of gas particles as temperature increases from 200 K to 450 K.

3A.2 Boyle's Law

Boyle's Law describes the inverse relationship between pressure and volume at constant temperature.

• Boyle's Law: $P_1V_1 = P_2V_2$

• Example: What will be the volume of a gas at 2.0 L and 0.8 atm if the pressure is increased to 1.6 atm?

3A.3 Charles's Law

Charles's Law describes the direct relationship between volume and temperature at constant pressure.

• Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$

• Example: A balloon has a volume of 1.5 L at 300 K. What is its volume at 450 K at constant pressure?

3A.4 Gay-Lussac's Law

Gay-Lussac's Law describes the direct relationship between pressure and temperature at constant volume.

• Gay-Lussac's Law: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$

• Example: A sealed container has a pressure of 1.2 atm at 250 K. What is the pressure at 400 K?

3A.5 Combined Gas Law

The combined gas law relates pressure, volume, and temperature for a fixed amount of gas.

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

• Example: A gas occupies 3.0 L at 2.0 atm and 300 K. What volume will it occupy at 1.0 atm and 450 K?

3A.6 Ideal Gas Law

The ideal gas law relates pressure, volume, temperature, and moles of gas.

• Ideal Gas Law: $PV = nRT$

• Example: How many moles of gas are present in a 10.0 L container at 2.0 atm and 300 K?

3A.7 Kinetic Molecular Theory & Deviations

Kinetic molecular theory explains the behavior of gases and deviations from ideality.

• Real Gases: Deviate from ideal behavior at high pressure and low temperature.

• Example: Under what conditions does a real gas behave least like an ideal gas?

Unit 6: Thermochemistry

6.1 Endothermic and Exothermic Processes

Thermochemistry studies energy changes in chemical reactions.

• Endothermic Process: Absorbs energy from surroundings.

• Exothermic Process: Releases energy to surroundings.

• Example: Is dissolving ammonium nitrate in water endothermic or exothermic?

6.3 Heat Transfer and Calorimetry

Calorimetry measures heat transfer in physical and chemical processes.

• Heat Transfer Equation: $q = m c \Delta T$

• Example: Calculate the heat absorbed by 50.0 g of water as its temperature increases from 20.0°C to 30.0°C ($c = 4.18$ J/g·°C).

6.4 Reaction Enthalpy (ΔH)

Reaction enthalpy is the heat change at constant pressure.

• ΔH: Positive for endothermic, negative for exothermic reactions.

• Example: If a reaction releases 125 kJ of energy, what is the sign and magnitude of ΔH?

6.5 Hess's Law

Hess's Law allows calculation of enthalpy changes for reactions using known enthalpies of related reactions.

• Hess's Law: $\Delta H_{rxn} = \sum \Delta H_{products} - \sum \Delta H_{reactants}$

• Example: Given reactions, calculate ΔH for a target reaction.

6.6 Bond Enthalpy

Bond enthalpy is the energy required to break a bond in a molecule.

• Bond Enthalpy: Breaking bonds requires energy; forming bonds releases energy.

• Example: Is a reaction that forms stronger bonds than it breaks endothermic or exothermic?

Unit 4: Chemical Reactions

4.1 Introduction to Chemical Reactions

Chemical reactions involve the transformation of substances via breaking and forming of bonds.

• Types of Reactions: Precipitation, acid-base, redox, combustion, single replacement.

• Example: Predict the products of the reaction between aqueous silver nitrate and sodium chloride.

4.2 Net Ionic Equations

Net ionic equations show only the species that participate in the reaction.

• Net Ionic Equation: Remove spectator ions to focus on the actual chemical change.

• Example: Write the net ionic equation for the reaction of HCl(aq) with NaOH(aq).

4.3 Stoichiometry (including limiting and excess reactants)

Stoichiometry uses balanced equations to relate quantities of reactants and products.

• Stoichiometric Calculations: Use mole ratios from balanced equations.

• Example: How many grams of CO2 are produced when 3.00 g of C4H10 reacts completely with excess O2?

4.4 Titrations and Molarity Calculations

Titrations determine unknown concentrations using a standard solution.

• Molarity: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Example: If 20.0 mL of NaOH titrates 20.0 mL of HCl, what is the molarity of HCl?

4.5 Oxidation-Reduction (Redox)

Redox reactions involve the transfer of electrons between species.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Example: Identify oxidation states and write net equations for redox reactions.

4.6 Balancing Reactions

Balancing chemical equations ensures the conservation of mass and charge.

• Balancing: Adjust coefficients to equalize the number of atoms of each element on both sides.

• Example: Balance the reaction: $\text{C}_4\text{H}_{10} + \text{O}_2 \rightarrow \text{CO}_2 + \text{H}_2\text{O}$

Additional info: These notes are based on an AP Chemistry midterm objective sheet and cover foundational topics in general chemistry, including atomic structure, bonding, gases, thermochemistry, and chemical reactions. Practice problems and objectives are included to guide study and application.