Exam 3 Practice Problems – Step-by-Step Guidance

Q1. What is the formula of iron(III) bromate?

Background

Topic: Naming and writing formulas for ionic compounds, specifically transition metal compounds with polyatomic ions.

This question tests your ability to write the correct chemical formula for a compound containing a transition metal with a specified oxidation state and a polyatomic ion.

Key Terms and Formulas

• Iron(III): The Roman numeral III indicates a +3 charge on the iron ion ($\mathrm{Fe^{3+}}$).

• Bromate: The polyatomic ion bromate has the formula $\mathrm{BrO_3^-}$.

Step-by-Step Guidance

1. Identify the charges of the ions: Iron(III) is $\mathrm{Fe^{3+}}$ and bromate is $\mathrm{BrO_3^-}$.

2. Determine the ratio of ions needed to balance the charges so the compound is neutral.

3. Write the formula by combining the ions in the correct ratio, using parentheses if more than one polyatomic ion is needed.

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Q2. What is the VSEPR geometry of sulfur dichloride (SCl2)?

Background

Topic: Molecular geometry using VSEPR theory.

This question tests your ability to predict the shape of a molecule based on the number of bonding pairs and lone pairs around the central atom.

Key Terms and Formulas

• VSEPR Theory: Valence Shell Electron Pair Repulsion theory is used to predict the geometry of molecules.

• Lone pairs: Non-bonding pairs of electrons on the central atom affect the shape.

Step-by-Step Guidance

1. Draw the Lewis structure for SCl2 to determine the number of bonding pairs and lone pairs on the central atom (sulfur).

2. Count the total electron domains (bonding pairs + lone pairs) around sulfur.

3. Use the VSEPR chart to determine the molecular geometry based on the number of electron domains and lone pairs.

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Q3. Draw the Lewis structure for BF3 and identify the number of lone pairs on the central atom.

Background

Topic: Lewis structures and electron counting.

This question tests your ability to draw a Lewis structure and count the number of lone pairs on the central atom (boron).

Key Terms and Formulas

• Lewis structure: A diagram showing the bonding between atoms and the lone pairs of electrons in a molecule.

• Lone pair: A pair of valence electrons not shared with another atom.

Step-by-Step Guidance

1. Count the total number of valence electrons for BF3 (boron and fluorine).

2. Draw the skeletal structure with boron as the central atom and attach three fluorine atoms.

3. Distribute the remaining electrons to satisfy the octet rule for the fluorine atoms first, then check boron's octet.

4. Count the number of lone pairs on the central atom (boron).

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Q4. Is nitrogen more electronegative than arsenic?

Background

Topic: Periodic trends – electronegativity.

This question tests your understanding of how electronegativity changes across the periodic table.

Key Terms and Formulas

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

• Electronegativity generally increases across a period (left to right) and decreases down a group (top to bottom).

Step-by-Step Guidance

1. Locate nitrogen and arsenic on the periodic table.

2. Recall the trend for electronegativity across periods and down groups.

3. Compare the relative positions of nitrogen and arsenic to determine which is more electronegative.

Try solving on your own before revealing the answer!

Q5. What is the bond angle for SCl2?

Background

Topic: Molecular geometry and bond angles.

This question tests your ability to relate molecular shape to bond angles, especially when lone pairs are present.

Key Terms and Formulas

• Bent geometry: The presence of lone pairs on the central atom reduces the bond angle from the ideal tetrahedral angle.

• Tetrahedral angle: $109.5^\circ$ is the ideal angle for four electron domains.

Step-by-Step Guidance

1. Determine the electron geometry of SCl2 using VSEPR theory.

2. Recall the ideal bond angle for a tetrahedral arrangement.

3. Consider how the presence of lone pairs affects the bond angle compared to the ideal value.

Try solving on your own before revealing the answer!

Q6. Identify the polyatomic ion: SO42-

Background

Topic: Polyatomic ions and their names.

This question tests your ability to recognize and name common polyatomic ions.

Key Terms and Formulas

• SO42-: A common polyatomic ion containing sulfur and oxygen.

Step-by-Step Guidance

1. Recall the names and formulas of common polyatomic ions containing sulfur and oxygen.

2. Match the given formula to its correct name.

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Q7. Which of the following species is best described by three equivalent resonance structures?

Background

Topic: Resonance structures in molecules and ions.

This question tests your understanding of resonance and the ability to identify species with multiple equivalent resonance forms.

Key Terms and Formulas

• Resonance: The concept that some molecules can be represented by two or more valid Lewis structures.

• Equivalent resonance structures: Structures that differ only in the placement of electrons, not atoms.

Step-by-Step Guidance

1. Draw the Lewis structures for each species listed in the options.

2. Determine if the species can be represented by three equivalent resonance structures.

3. Look for delocalized electrons and symmetry in the molecule or ion.

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Q8. What is the formal charge of the least electronegative element in sulfur dioxide (SO2)?

Background

Topic: Formal charge calculation in Lewis structures.

This question tests your ability to assign formal charges to atoms in a molecule, especially focusing on the least electronegative atom.

Key Terms and Formulas

• Formal charge formula:

  $\text{Formal charge} = \text{valence electrons} - (\text{non-bonding electrons} + \frac{1}{2} \times \text{bonding electrons})$

• Electronegativity: The least electronegative element is usually the central atom.

Step-by-Step Guidance

1. Draw the Lewis structure for SO2.

2. Identify the least electronegative atom (sulfur or oxygen).

3. Count the valence electrons, non-bonding electrons, and bonding electrons for that atom.

4. Apply the formal charge formula to calculate the formal charge for the least electronegative atom.

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Q9. How many valence electrons do group 16 (6A) elements have?

Background

Topic: Periodic table and valence electrons.

This question tests your knowledge of how to determine the number of valence electrons for main group elements based on their group number.

Key Terms and Formulas

• Valence electrons: Electrons in the outermost shell of an atom, important for bonding.

• For main group elements, the group number (for groups 1A–8A) indicates the number of valence electrons.

Step-by-Step Guidance

1. Locate group 16 (6A) on the periodic table.

2. Recall the rule for determining valence electrons from the group number.

3. State the number of valence electrons for elements in this group.

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Q10. What is the charge on ions formed by group 2 (2A) elements?

Background

Topic: Periodic trends and ion formation.

This question tests your understanding of how main group elements form ions and what charges they typically have.

Key Terms and Formulas

• Group 2 (2A) elements: Also known as alkaline earth metals.

• These elements lose electrons to achieve a noble gas configuration.

Step-by-Step Guidance

1. Identify the number of valence electrons in group 2 elements.

2. Determine how many electrons they lose to achieve a stable configuration.

3. Assign the resulting charge to the ion formed.

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Q11. Determine the polarity of NH3 (ammonia).

Background

Topic: Molecular polarity.

This question tests your ability to determine if a molecule is polar or nonpolar based on its shape and the electronegativity of its atoms.

Key Terms and Formulas

• Polarity: A molecule is polar if it has a net dipole moment due to differences in electronegativity and molecular geometry.

• NH3: Has a trigonal pyramidal shape with a lone pair on nitrogen.

Step-by-Step Guidance

1. Draw the Lewis structure for NH3 and determine its molecular geometry.

2. Consider the electronegativity difference between nitrogen and hydrogen.

3. Assess whether the molecular shape allows for a net dipole moment.

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Q12. What is the chemical formula for iron(II) nitrate?

Background

Topic: Writing formulas for ionic compounds with transition metals and polyatomic ions.

This question tests your ability to write the correct formula for a compound containing a transition metal with a specified oxidation state and a polyatomic ion.

Key Terms and Formulas

• Iron(II): The Roman numeral II indicates a +2 charge on the iron ion ($\mathrm{Fe^{2+}}$).

• Nitrate: The polyatomic ion nitrate has the formula $\mathrm{NO_3^-}$.

Step-by-Step Guidance

1. Identify the charges of the ions: Iron(II) is $\mathrm{Fe^{2+}}$ and nitrate is $\mathrm{NO_3^-}$.

2. Determine the ratio of ions needed to balance the charges so the compound is neutral.

3. Write the formula by combining the ions in the correct ratio, using parentheses if more than one polyatomic ion is needed.

Try solving on your own before revealing the answer!