Exam 3 Practice Problems – Step-by-Step Guidance

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

Background

Topic: Naming and Writing Formulas for Ionic Compounds

This question tests your ability to write the chemical formula for an ionic compound given the name, specifically for a transition metal with 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 bromate ion is a polyatomic ion with the formula $\mathrm{BrO_3^-}$.

Step-by-Step Guidance;800357883=0115?

1. Write the symbols and charges for each ion: $\mathrm{Fe^{3+}}$ and $\mathrm{BrO_3^-}$.

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

3. Use parentheses around the polyatomic ion if more than one is needed.

4. Write the final formula by combining the ions in the correct ratio.

Try solving on your own before revealing the answer!

Final Answer: Fe(BrO3)3

Iron(III) has a +3 charge and bromate has a -1 charge, so three bromate ions are needed to balance one iron(III) ion.

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

Background

Topic: Molecular Geometry (VSEPR Theory)

This question tests your understanding of how electron pairs around a central atom determine the shape of a molecule.

Key Terms and Formulas

• VSEPR Theory: Valence Shell Electron Pair Repulsion theory predicts the geometry of molecules based on repulsion between electron pairs.

• SCl2: Sulfur is the central atom with two chlorine atoms bonded to it.

Step-by-Step Guidance

1. Count the total number of valence electrons for SCl2.

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

3. Use VSEPR notation (AXnEm) to classify the geometry based on the number of bonding and lone pairs.

4. Identify the molecular geometry corresponding to this arrangement.

Try solving on your own before revealing the answer!

Final Answer: Bent

Sulfur dichloride has two bonding pairs and two lone pairs on the central atom, resulting in a bent geometry.

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

Background

Topic: Lewis Structures and Electron Pair Counting

This question tests your ability to draw Lewis structures and count lone pairs on the central atom.

Key Terms and Formulas

• Lewis Structure: A diagram showing the arrangement of valence electrons among atoms 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 in BF3 (boron and fluorine).

2. Draw single bonds between boron and each fluorine atom.

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

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

Try solving on your own before revealing the answer!

Final Answer: None

Boron in BF3 has no lone pairs; all its valence electrons are involved in bonding with fluorine atoms.

Q4. Is nitrogen more electronegative than arsenic?

Background

Topic: Electronegativity Trends

This question tests your understanding of periodic trends, specifically electronegativity, across and down the periodic table.

Key Terms and Formulas

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

• Electronegativity 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 (both are in group 15, but different periods).

2. Recall the trend: electronegativity increases as you move up and to the right.

3. Compare their positions to determine which is more electronegative.

Try solving on your own before revealing the answer!

Final Answer: True

Nitrogen is above arsenic in the same group, so it is more electronegative.

Q5. What is the bond angle for SCl2?

Background

Topic: Molecular Geometry and Bond Angles

This question tests your ability to relate molecular geometry to bond angles using VSEPR theory.

Key Terms and Formulas

• Bent Geometry: Molecules with two bonding pairs and two lone pairs (like SCl2) have bond angles less than the ideal tetrahedral angle.

• Tetrahedral Angle: $109.5^\circ$

Step-by-Step Guidance

1. Recall the electron geometry for SCl2 (tetrahedral electron geometry, bent molecular geometry).

2. Know that lone pairs compress bond angles below the ideal tetrahedral angle.

3. Compare the bond angle to standard values (109.5°, 120°, 180°) and determine if it is less than these values.

Try solving on your own before revealing the answer!

Final Answer: Less than 109.5°

Lone pairs on sulfur reduce the bond angle below the ideal tetrahedral angle of 109.5°.

Q6. Identify the polyatomic ion SO42-.

Background

Topic: Polyatomic Ions

This question tests your ability to recognize and name common polyatomic ions based on their formulas.

Key Terms and Formulas

• SO42-: A polyatomic ion consisting of one sulfur atom and four oxygen atoms with a -2 charge.

Step-by-Step Guidance

1. Recall the names and formulas of common polyatomic ions.

2. Match SO42- to its correct name from the list provided.

Try solving on your own before revealing the answer!

Final Answer: Sulfate

SO42- is the formula for the sulfate ion.

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

Background

Topic: Resonance Structures

This question tests your understanding of resonance and the ability to identify molecules or ions that have multiple equivalent resonance forms.

Key Terms and Formulas

• Resonance Structures: Different Lewis structures for the same molecule or ion that show delocalization of electrons.

• Equivalent resonance structures have the same arrangement of atoms but different arrangements of electrons.

Step-by-Step Guidance

1. Recall which molecules or ions commonly have three equivalent resonance structures (often those with three identical terminal atoms bonded to a central atom).

2. Consider the Lewis structures for each option and see if the double bonds or negative charges can be delocalized in three ways.

Try solving on your own before revealing the answer!

Final Answer: SO3

SO3 has three equivalent resonance structures, each with a double bond to a different oxygen atom.

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

Background

Topic: Formal Charge Calculation

This question tests your ability to assign formal charges to atoms in a molecule using the Lewis structure.

Key Terms and Formulas

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

• Electronegativity: Oxygen is more electronegative than sulfur.

Step-by-Step Guidance

1. Draw the Lewis structure for SO2.

2. Identify the least electronegative atom (sulfur).

3. Count the number of valence electrons, non-bonding electrons, and bonding electrons for sulfur.

4. Apply the formal charge formula to sulfur.

Try solving on your own before revealing the answer!

Final Answer: 0

The formal charge on sulfur in SO2 is zero when the most stable Lewis structure is drawn.

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 chemical bonding.

• Group 16 (6A) elements include oxygen, sulfur, selenium, etc.

Step-by-Step Guidance

1. Recall that the group number for main group elements indicates the number of valence electrons (for groups 1A–8A).

2. Group 16 (6A) elements are in the sixth column of the main group elements.

3. Assign the correct number of valence electrons based on the group number.

Try solving on your own before revealing the answer!

Final Answer: 6

Group 16 (6A) elements have 6 valence electrons.