General Chemistry Study Guide: Chapters 9–11 (Electron Configurations, Bonding, Molecular Shapes)

Q1. Write electron configurations for elements. Understand Pauli exclusion principle and Hund’s rule (9.3)

Background

Topic: Electron Configuration

This question tests your understanding of how electrons are arranged in atoms, including the rules that govern their arrangement: the Pauli exclusion principle and Hund’s rule.

Key Terms and Formulas

• Electron configuration: The distribution of electrons among the orbitals of an atom.

• Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers.

• Hund’s rule: Electrons fill degenerate orbitals singly first, with parallel spins, before pairing up.

Step-by-Step Guidance

1. Identify the atomic number of the element to determine the total number of electrons.

2. List the order of orbital filling using the Aufbau principle (e.g., 1s, 2s, 2p, 3s, 3p, etc.).

3. Apply the Pauli exclusion principle: Each orbital can hold a maximum of two electrons with opposite spins.

4. Apply Hund’s rule: When filling orbitals of equal energy (like the three 2p orbitals), place one electron in each before pairing.

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Q2. Draw orbital diagrams for elements (9.3)

Background

Topic: Orbital Diagrams

This question tests your ability to visually represent electron configurations using boxes and arrows to show electron spins.

Key Terms and Formulas

• Orbital diagram: A pictorial representation of electron configuration, showing orbitals as boxes and electrons as arrows.

• Spin: Represented by up or down arrows.

Step-by-Step Guidance

1. Determine the electron configuration for the element.

2. Draw boxes for each orbital (e.g., one for 1s, three for 2p).

3. Fill the boxes with arrows, following Hund’s rule and the Pauli exclusion principle.

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Q3. Write electron configurations based on periodic table location. Understand core and valence electrons (9.4)

Background

Topic: Periodic Table and Electron Configuration

This question tests your ability to use the periodic table to determine electron configurations and distinguish between core and valence electrons.

Key Terms and Formulas

• Core electrons: Electrons in inner shells, not involved in bonding.

• Valence electrons: Electrons in the outermost shell, involved in bonding.

Step-by-Step Guidance

1. Locate the element on the periodic table to determine its group and period.

2. Write the electron configuration, using noble gas notation for core electrons.

3. Identify the valence electrons based on the highest principal quantum number.

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Q4. Predict effective nuclear charge and relative atomic sizes of elements (9.6)

Background

Topic: Effective Nuclear Charge and Atomic Size

This question tests your understanding of how the effective nuclear charge ($Z_{\text{eff}}$) affects atomic size across periods and groups.

Key Terms and Formulas

• Effective nuclear charge ($Z_{\text{eff}}$): The net positive charge experienced by valence electrons.

• Atomic size: The radius of an atom, influenced by $Z_{\text{eff}}$.

• $Z_{\text{eff}} = Z - S$

• Where $Z$ is the atomic number and $S$ is the shielding constant.

Step-by-Step Guidance

1. Identify the atomic number ($Z$) for the element.

2. Estimate the shielding constant ($S$) based on the number of core electrons.

3. Calculate $Z_{\text{eff}}$ using $Z_{\text{eff}} = Z - S$.

4. Compare $Z_{\text{eff}}$ and atomic size for elements across a period and down a group.

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Q5. Analyze ions in terms of magnetic properties (9.7)

Background

Topic: Magnetism in Ions

This question tests your understanding of paramagnetism and diamagnetism based on electron configurations.

Key Terms and Formulas

• Paramagnetic: Ions with unpaired electrons, attracted to magnetic fields.

• Diamagnetic: Ions with all electrons paired, not attracted to magnetic fields.

Step-by-Step Guidance

1. Determine the electron configuration for the ion.

2. Identify if there are any unpaired electrons.

3. Classify the ion as paramagnetic or diamagnetic based on the presence of unpaired electrons.

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Q6. Predict the relative size of ions from periodic trends (9.7)

Background

Topic: Ionic Size and Periodic Trends

This question tests your ability to predict how ionic size changes with charge and position on the periodic table.

Key Terms and Formulas

• Cation: Positively charged ion, usually smaller than the parent atom.

• Anion: Negatively charged ion, usually larger than the parent atom.

Step-by-Step Guidance

1. Identify the charge and electron configuration of the ion.

2. Compare the ionic radius to the atomic radius.

3. Use periodic trends to predict relative sizes among ions.

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Q7. Predict relative ionization energies for atoms and ions based on periodic trends (9.7)

Background

Topic: Ionization Energy and Periodic Trends

This question tests your understanding of how ionization energy changes across periods and groups.

Key Terms and Formulas

• Ionization energy: The energy required to remove an electron from an atom or ion.

• Periodic trend: Ionization energy increases across a period and decreases down a group.

Step-by-Step Guidance

1. Identify the position of the atom or ion on the periodic table.

2. Recall the general trend for ionization energy across periods and down groups.

3. Consider electron configuration and effective nuclear charge.

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Q8. Write Lewis symbols for main-group elements (10.3)

Background

Topic: Lewis Symbols

This question tests your ability to represent valence electrons for main-group elements using Lewis symbols.

Key Terms and Formulas

• Lewis symbol: A diagram showing the element’s symbol and dots for each valence electron.

Step-by-Step Guidance

1. Determine the number of valence electrons for the element.

2. Draw the element’s symbol.

3. Place dots around the symbol to represent valence electrons.

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Q9. Write and use Lewis symbols to describe and analyze ionic compounds (10.4)

Background

Topic: Lewis Symbols and Ionic Compounds

This question tests your ability to use Lewis symbols to show electron transfer and formation of ions in ionic compounds.

Key Terms and Formulas

• Ionic compound: Formed by transfer of electrons from metal to nonmetal.

• Lewis symbol: Shows electron transfer and resulting ions.

Step-by-Step Guidance

1. Draw Lewis symbols for the metal and nonmetal.

2. Show the transfer of electrons from the metal to the nonmetal.

3. Represent the resulting ions with brackets and charges.

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Q10. Compare the relative lattice energies of different compounds (10.4)

Background

Topic: Lattice Energy

This question tests your understanding of how lattice energy varies with ionic charge and size.

Key Terms and Formulas

• Lattice energy: The energy released when ions form a solid lattice.

• Depends on charge and size of ions.

• $U = k \frac{Q_1 Q_2}{r}$

• Where $U$ is lattice energy, $Q_1$ and $Q_2$ are ion charges, $r$ is distance between ions, $k$ is a constant.

Step-by-Step Guidance

1. Identify the charges of the ions in each compound.

2. Compare the sizes (radii) of the ions.

3. Use the formula to predict which compound has higher lattice energy.

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Q11. Calculate lattice energies of ionic solids (10.4)

Background

Topic: Lattice Energy Calculation

This question tests your ability to use the lattice energy formula to calculate the energy for ionic solids.

Key Terms and Formulas

• Lattice energy formula: $U = k \frac{Q_1 Q_2}{r}$

Step-by-Step Guidance

1. Identify the charges ($Q_1$, $Q_2$) and radii ($r$) for the ions.

2. Plug values into the formula.

3. Calculate the intermediate steps, stopping before the final calculation.

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Q12. Classify bonds as pure covalent, polar covalent, or ionic (10.5–10.6)

Background

Topic: Bond Classification

This question tests your ability to classify bonds based on electronegativity differences.

Key Terms and Formulas

• Electronegativity: The ability of an atom to attract electrons.

• Bond types: Pure covalent (equal sharing), polar covalent (unequal sharing), ionic (transfer).

Step-by-Step Guidance

1. Determine the electronegativity values for the atoms involved.

2. Calculate the difference in electronegativity.

3. Classify the bond based on the difference: small (covalent), moderate (polar covalent), large (ionic).

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Q13. Write and use Lewis structures to describe and analyze molecular compounds and polyatomic ions (10.7)

Background

Topic: Lewis Structures for Molecules and Polyatomic Ions

This question tests your ability to draw and interpret Lewis structures for molecules and ions.

Key Terms and Formulas

• Lewis structure: Shows how atoms are bonded and the arrangement of valence electrons.

Step-by-Step Guidance

1. Count the total number of valence electrons for all atoms.

2. Arrange atoms and connect with single bonds.

3. Distribute remaining electrons to satisfy the octet rule.

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Q14. Write the best possible Lewis structure using formal charge (10.8)

Background

Topic: Formal Charge and Lewis Structures

This question tests your ability to use formal charge to determine the most stable Lewis structure.

Key Terms and Formulas

• Formal charge formula: $\text{Formal charge} = \text{Valence electrons} - (\text{Nonbonding electrons} + \frac{1}{2} \text{Bonding electrons})$

Step-by-Step Guidance

1. Draw possible Lewis structures for the molecule or ion.

2. Calculate formal charges for each atom in each structure.

3. Choose the structure with the lowest formal charges and most negative charge on the most electronegative atom.

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Q15. Calculate formal charge (10.8)

Background

Topic: Formal Charge Calculation

This question tests your ability to calculate formal charge for atoms in a molecule or ion.

Key Terms and Formulas

• Formal charge formula: $\text{Formal charge} = \text{Valence electrons} - (\text{Nonbonding electrons} + \frac{1}{2} \text{Bonding electrons})$

Step-by-Step Guidance

1. Identify the number of valence electrons for the atom.

2. Count the number of nonbonding electrons and bonding electrons.

3. Plug values into the formal charge formula.

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Q16. Write Lewis structures for molecules and ions that are exceptions to the octet rule (10.9)

Background

Topic: Exceptions to the Octet Rule

This question tests your ability to draw Lewis structures for molecules and ions that do not follow the octet rule (e.g., odd number of electrons, expanded octet).

Key Terms and Formulas

• Octet rule: Most atoms aim for eight electrons in their valence shell.

• Exceptions: Odd-electron molecules, expanded octet, incomplete octet.

Step-by-Step Guidance

1. Count total valence electrons.

2. Arrange atoms and distribute electrons, noting exceptions.

3. Check for expanded or incomplete octets as appropriate.

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Q17. Analyze bonds in terms of bond energy and bond length (10.10)

Background

Topic: Bond Energy and Bond Length

This question tests your understanding of how bond energy and bond length are related.

Key Terms and Formulas

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

• Bond length: The distance between nuclei of bonded atoms.

• Generally, shorter bonds are stronger (higher bond energy).

Step-by-Step Guidance

1. Identify the types of bonds (single, double, triple).

2. Compare bond lengths and bond energies for each type.

3. Relate bond order to bond energy and length.

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Q18. Calculate bond energy in a chemical reaction (10.10)

Background

Topic: Bond Energy Calculation

This question tests your ability to calculate the energy change in a reaction using bond energies.

Key Terms and Formulas

• Bond energy: Energy required to break a bond.

• Formula: $\Delta H = \sum \text{(bond energies of bonds broken)} - \sum \text{(bond energies of bonds formed)}$

Step-by-Step Guidance

1. List all bonds broken and formed in the reaction.

2. Find bond energies for each bond type.

3. Set up the calculation using the formula above.

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Q19. Predict the basic shapes of molecules according to VSEPR theory (11.2)

Background

Topic: VSEPR Theory and Molecular Shape

This question tests your ability to use VSEPR theory to predict molecular shapes based on electron group arrangement.

Key Terms and Formulas

• VSEPR theory: Valence Shell Electron Pair Repulsion theory.

• Electron groups: Bonding pairs and lone pairs around the central atom.

Step-by-Step Guidance

1. Count the number of electron groups around the central atom.

2. Use VSEPR theory to determine the basic shape (linear, trigonal planar, tetrahedral, etc.).

3. Draw or visualize the shape based on electron group arrangement.

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Q20. Predict the basic shapes of molecules according to VSEPR theory. Predict how lone pairs and electron groups affect molecular geometry (11.3–11.4)

Background

Topic: VSEPR Theory and Molecular Geometry

This question tests your ability to predict how lone pairs and electron groups influence the geometry of molecules.

Key Terms and Formulas

• Lone pairs: Nonbonding electron pairs on the central atom.

• Molecular geometry: The actual shape of the molecule, considering lone pairs.

Step-by-Step Guidance

1. Count the number of bonding pairs and lone pairs on the central atom.

2. Determine the electron group geometry using VSEPR theory.

3. Adjust the predicted shape to account for lone pairs, which can compress bond angles.

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