General Chemistry Study Guide: Periodic Properties, Bonding, and Molecular Structure

Q1. Choose the ground state electron configuration for Ti2+.

Background

Topic: Electron Configuration of Transition Metal Ions

This question tests your understanding of how electrons are removed from transition metals to form cations, and how to write condensed electron configurations using the noble gas core notation.

Key Terms and Formulas:

• Electron Configuration: The arrangement of electrons in an atom or ion.

• Transition Metals: For cations, electrons are removed first from the 4s orbital, then from the 3d orbital.

Step-by-Step Guidance

1. Write the ground state electron configuration for neutral Ti (atomic number 22): $\mathrm{Ti}: [\mathrm{Ar}]\,4s^2\,3d^2$.

2. When forming a cation, remove electrons first from the 4s orbital, then from the 3d orbital.

3. Remove two electrons to form Ti2+. Consider which orbitals lose electrons first.

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Q2. Choose the ground state electron configuration for Cr3+.

Background

Topic: Electron Configuration of Transition Metal Ions

This question tests your ability to determine the electron configuration for a transition metal cation, considering the special stability of half-filled and fully-filled d subshells.

Key Terms and Formulas:

• Electron Configuration: The arrangement of electrons in an atom or ion.

• Cr (Chromium): Has an unusual configuration due to stability of half-filled d subshells.

Step-by-Step Guidance

1. Write the ground state electron configuration for neutral Cr (atomic number 24): $\mathrm{Cr}: [\mathrm{Ar}]\,4s^1\,3d^5$.

2. Remove three electrons to form Cr3+. Remove from 4s first, then 3d.

3. Write the resulting configuration after removing the electrons.

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Q3. Place the following elements in order of decreasing atomic radius: Cd, Rb, Xe.

Background

Topic: Periodic Trends – Atomic Radius

This question tests your understanding of how atomic radius changes across periods and down groups in the periodic table.

Key Terms and Concepts:

• Atomic Radius: The size of an atom, typically increases down a group and decreases across a period.

• Periodic Trend: Atomic radius decreases left to right, increases top to bottom.

Step-by-Step Guidance

1. Locate Cd, Rb, and Xe on the periodic table.

2. Recall that atomic radius increases down a group and decreases across a period.

3. Arrange the elements from largest to smallest atomic radius based on their positions.

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Q4. Place the following in order of increasing radius: Se2−, Mg2+, Na+.

Background

Topic: Ionic Radii and Isoelectronic Series

This question tests your understanding of how ionic size changes with charge and across isoelectronic ions (ions with the same number of electrons).

Key Terms and Concepts:

• Ionic Radius: The size of an ion.

• Isoelectronic Series: Ions with the same number of electrons but different nuclear charges.

Step-by-Step Guidance

1. Determine the number of electrons in each ion.

2. Recall that for isoelectronic ions, the more positive the charge, the smaller the radius.

3. Arrange the ions in order of increasing radius.

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Q5. Place the following in order of increasing first ionization energy (IE1): N, F, As.

Background

Topic: Ionization Energy Trends

This question tests your understanding of periodic trends in ionization energy, which is the energy required to remove the outermost electron from a neutral atom.

Key Terms and Concepts:

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

Step-by-Step Guidance

1. Locate N, F, and As on the periodic table.

2. Recall the trend: IE increases left to right and decreases top to bottom.

3. Arrange the elements in order of increasing IE1.

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Q6. Which reaction below represents the electron affinity of K?

Background

Topic: Electron Affinity

This question tests your understanding of electron affinity, which is the energy change when an atom in the gas phase gains an electron.

Key Terms and Concepts:

• Electron Affinity: The process of adding an electron to a neutral atom in the gas phase to form an anion.

Step-by-Step Guidance

1. Recall the definition of electron affinity.

2. Identify which reaction shows a neutral K atom in the gas phase gaining an electron to form K−(g).

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Q7. What period 3 element has the following ionization energies (all in kJ/mol)?

Background

Topic: Successive Ionization Energies

This question tests your ability to interpret patterns in ionization energies to identify an element based on large jumps between successive ionization energies.

Key Terms and Concepts:

• Successive Ionization Energies: Large jumps indicate removal of electrons from a new shell (core electrons).

Step-by-Step Guidance

1. Look for the largest jump in ionization energy values.

2. Determine how many valence electrons the element has based on where the jump occurs.

3. Match the number of valence electrons to a period 3 element.

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Q8. Place the following in order of decreasing magnitude of lattice energy: K2O, Rb2S, Li2O.

Background

Topic: Lattice Energy

This question tests your understanding of factors affecting lattice energy, including ionic charge and ionic radius.

Key Terms and Formulas:

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

• Key Formula: $U \propto \frac{Q_1 Q_2}{r}$ (where $Q_1$ and $Q_2$ are the charges, $r$ is the distance between ions)

Step-by-Step Guidance

1. Compare the charges on the ions in each compound (all are +1 and -2).

2. Compare the sizes of the ions (smaller ions = higher lattice energy).

3. Arrange the compounds in order of decreasing lattice energy.

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Q9. Identify the longest bond.

Background

Topic: Bond Order and Bond Length

This question tests your understanding of the relationship between bond order (single, double, triple) and bond length.

Key Terms and Concepts:

• Bond Order: Single bond (1), double bond (2), triple bond (3).

• Bond Length: Inversely related to bond order; single bonds are longest, triple bonds are shortest.

Step-by-Step Guidance

1. Recall the trend: single bond > double bond > triple bond in terms of length.

2. Identify which bond type is the longest.

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Q10. Define electronegativity.

Background

Topic: Electronegativity

This question tests your understanding of the definition of electronegativity and its role in chemical bonding.

Key Terms and Concepts:

• Electronegativity: The ability of an atom to attract electrons in a chemical bond.

Step-by-Step Guidance

1. Recall the definition of electronegativity.

2. Identify the correct choice that matches this definition.

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Q11. Choose the best Lewis structure for BeF2.

Background

Topic: Lewis Structures

This question tests your ability to draw and evaluate Lewis structures for simple molecules, considering the octet rule and formal charges.

Key Terms and Concepts:

• Lewis Structure: A diagram showing the arrangement of valence electrons among atoms in a molecule.

• Octet Rule: Atoms tend to form bonds to have eight electrons in their valence shell (with exceptions).

Step-by-Step Guidance

1. Count the total number of valence electrons for BeF2.

2. Draw possible Lewis structures, ensuring Be and F have appropriate numbers of electrons.

3. Check for formal charges and octet rule satisfaction.

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Q12. Give the number of valence electrons for IF5.

Background

Topic: Counting Valence Electrons

This question tests your ability to determine the total number of valence electrons in a molecule, which is essential for drawing Lewis structures.

Key Terms and Concepts:

• Valence Electrons: Electrons in the outermost shell of an atom, involved in bonding.

Step-by-Step Guidance

1. Find the number of valence electrons for Iodine (I) and Fluorine (F) from the periodic table.

2. Multiply the number of F atoms by the number of valence electrons for F.

3. Add the valence electrons from I and F to get the total.

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Q13. Choose the best Lewis structure for ICl5.

Background

Topic: Lewis Structures for Hypervalent Molecules

This question tests your ability to draw Lewis structures for molecules where the central atom can have more than eight electrons (expanded octet).

Key Terms and Concepts:

• Lewis Structure: Shows bonding and lone pairs in a molecule.

• Expanded Octet: Atoms in period 3 or higher can have more than 8 electrons.

Step-by-Step Guidance

1. Count the total number of valence electrons for ICl5.

2. Draw possible Lewis structures, ensuring Iodine can have more than 8 electrons.

3. Check for formal charges and octet rule satisfaction for Cl atoms.

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Q14. Draw the Lewis structure for CO32− including any valid resonance structures. Describe one resonance structure of the carbonate ion.

Background

Topic: Resonance Structures

This question tests your understanding of resonance in polyatomic ions and how to describe the bonding in terms of single and double bonds.

Key Terms and Concepts:

• Resonance Structures: Different valid Lewis structures for the same molecule, showing delocalization of electrons.

• Bond Order: Average number of bonds between atoms in resonance structures.

Step-by-Step Guidance

1. Count the total number of valence electrons for CO32−.

2. Draw all possible resonance structures, distributing electrons to satisfy the octet rule.

3. Describe one resonance structure in terms of single and double bonds.

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Q15. Use the bond energies provided to estimate ΔH°rxn for the reaction: 2 Br2(l) + C2H2(g) → C2H2Br4(l)

Background

Topic: Bond Energies and Enthalpy Calculations

This question tests your ability to use bond energies to estimate the enthalpy change of a reaction by considering bonds broken and formed.

Key Formula:

$\Delta H_{rxn} = \sum (\text{Bond energies of bonds broken}) - \sum (\text{Bond energies of bonds formed})$

Step-by-Step Guidance

1. List all bonds broken in the reactants and sum their bond energies.

2. List all bonds formed in the products and sum their bond energies.

3. Plug the values into the formula above to set up the calculation.

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Q16. Choose the bond below that is most polar.

Background

Topic: Bond Polarity and Electronegativity

This question tests your understanding of how differences in electronegativity between two atoms affect bond polarity.

Key Terms and Concepts:

• Bond Polarity: Determined by the difference in electronegativity between two atoms.

• Electronegativity: Use the Pauling scale to compare values.

Step-by-Step Guidance

1. Look up the electronegativity values for each pair of atoms.

2. Calculate the difference for each bond.

3. The bond with the largest difference is the most polar.

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Q17. In the best Lewis structure for XeF4, what is the formal charge on the F atom?

Background

Topic: Formal Charge Calculation

This question tests your ability to calculate formal charges in a Lewis structure, which helps determine the most stable structure.

Key Formula:

$\text{Formal Charge} = \text{Valence electrons} - (\text{Non-bonding electrons} + \frac{1}{2} \text{Bonding electrons})$

Step-by-Step Guidance

1. Draw the Lewis structure for XeF4.

2. Count the valence electrons, non-bonding electrons, and bonding electrons for F.

3. Plug these values into the formal charge formula for F.

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Q18. Give the molecular geometry and number of electron groups for BrF3.

Background

Topic: VSEPR Theory

This question tests your ability to use VSEPR theory to predict molecular geometry and count electron groups around the central atom.

Key Terms and Concepts:

• Electron Groups: Bonds and lone pairs around the central atom.

• Molecular Geometry: The shape determined by the positions of atoms (not lone pairs).

Step-by-Step Guidance

1. Draw the Lewis structure for BrF3.

2. Count the number of electron groups (bonds and lone pairs) around Br.

3. Use VSEPR theory to determine the molecular geometry.

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Q19. Identify the number of electron groups around a molecule with sp2 hybridization.

Background

Topic: Hybridization

This question tests your understanding of the relationship between hybridization and the number of electron groups around a central atom.

Key Terms and Concepts:

• sp2 Hybridization: Involves mixing one s and two p orbitals.

• Electron Groups: Each group can be a bond or a lone pair.

Step-by-Step Guidance

1. Recall the number of electron groups associated with sp2 hybridization.

2. Relate this to the geometry (trigonal planar).

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Q20. Give the approximate bond angle for a molecule with a trigonal planar shape.

Background

Topic: Molecular Geometry and Bond Angles

This question tests your knowledge of the bond angles associated with different molecular geometries.

Key Terms and Concepts:

• Trigonal Planar: Three electron groups around a central atom, all in one plane.

• Bond Angle: The angle between adjacent bonds.

Step-by-Step Guidance

1. Recall the ideal bond angle for a trigonal planar geometry.

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Q21. What is the molecular geometry of SCl4?

Background

Topic: VSEPR Theory and Molecular Geometry

This question tests your ability to predict the molecular geometry of a molecule with five electron groups (including lone pairs).

Key Terms and Concepts:

• VSEPR Theory: Predicts shapes based on repulsion between electron groups.

• Seesaw Geometry: Occurs with four bonds and one lone pair.

Step-by-Step Guidance

1. Draw the Lewis structure for SCl4.

2. Count the number of bonding pairs and lone pairs on S.

3. Use VSEPR to determine the geometry.

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Q22. Choose the compound below that contains at least one polar covalent bond, but is nonpolar.

Background

Topic: Molecular Polarity

This question tests your understanding of the difference between bond polarity and molecular polarity, and how molecular geometry affects overall polarity.

Key Terms and Concepts:

• Polar Covalent Bond: Bond between atoms with different electronegativities.

• Nonpolar Molecule: Molecule with symmetrical geometry so dipoles cancel.

Step-by-Step Guidance

1. Identify which compounds have polar bonds (difference in electronegativity).

2. Determine if the molecular geometry allows for dipole cancellation (nonpolar overall).

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Q23. Give the hybridization for the S in SCl6.

Background

Topic: Hybridization and Electron Geometry

This question tests your ability to determine the hybridization of the central atom based on the number of electron groups.

Key Terms and Concepts:

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals.

• sp3d2 Hybridization: Associated with six electron groups (octahedral geometry).

Step-by-Step Guidance

1. Draw the Lewis structure for SCl6.

2. Count the number of electron groups around S.

3. Match the number of groups to the correct hybridization scheme.

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Q24. List the number of sigma bonds and pi bonds in a double bond.

Background

Topic: Sigma and Pi Bonds

This question tests your understanding of the types of bonds present in single, double, and triple bonds.

Key Terms and Concepts:

• Sigma (σ) Bond: First bond formed between two atoms.

• Pi (π) Bond: Additional bonds in double/triple bonds.

Step-by-Step Guidance

1. Recall that a single bond is always a sigma bond.

2. In a double bond, one is sigma and the other is pi.

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Q25. Use MO theory to predict the bond order in He2+ ion. Is the He2+ bond a stronger or weaker bond than the He2 bond?

Background

Topic: Molecular Orbital Theory

This question tests your ability to use molecular orbital diagrams to calculate bond order and compare bond strengths.

Key Formula:

$\text{Bond Order} = \frac{1}{2}[(\text{number of bonding electrons}) - (\text{number of antibonding electrons})]$

Step-by-Step Guidance

1. Draw the MO diagram for He2+ and He2.

2. Count the number of electrons in bonding and antibonding orbitals for each species.

3. Calculate the bond order for He2+ and compare to He2.

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