Problem 67 - Chemistry: The Central Science in SI Units, Global Edition

Question

There are many Lewis structures you could draw for sulfuric acid, \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (each \(\mathrm{H}\) is bonded to an \(\mathrm{O}\) ). (a) What Lewis structure(s) would you draw to satisfy the octet rule? (b) What Lewis structure(s) would you draw to minimize formal charge?

Step-by-Step Solution

Answer

To satisfy the octet rule, the Lewis structure for H₂SO₄ would be: \( H-O-S-O \) | X X O O However, to minimize formal charges, the Lewis structure for H₂SO₄ would be: \( H-O-S(O_{3})(O_{4}) \) || | || O_{1} O_{2} In this structure, O₁ and O₂ have double bonds with sulfur, O₃ and O₄ have single bonds, and all atoms have formal charges of 0 or -1.

Step 1: Calculate valence electrons

H has 1, S has 6, and O has 6 valence electrons each. Since there are 2 H atoms, 1 S atom, and 4 O atoms, we will have a total of 2(1) + 6 + 4(6) = 2 + 6 + 24 = 32 valence electrons.

Step 2: Arrange atoms and add bonds

We will place the S atom in the center, with the 4 O atoms surrounding it and the two H atoms bonded to two of the O atoms. This gives us the following initial structure: \( H-O-S-O \) | X X O O Each S-O bond and O-H bond counts as 2 electrons. This accounts for 2(4) + 2(2) = 12 of the 32 valence electrons.

Step 3: Distribute remaining electrons

We have 32 (total) - 12 (used for bonds) = 20 electrons left. Each oxygen should have 6 electrons around it to fulfill the octet rule, so we will add 4 electrons to each of the two O atoms not bonded to H and 4 more electrons to each of the two O atoms bonded to H.

Step 4: Lewis structure satisfying the octet rule

The resulting Lewis structure for H₂SO₄ that satisfies the octet rule is: \( H-O-S-O \) | X X O O b. Drawing Lewis structure(s) to minimize formal charge#

Step 5: Calculate formal charges

We can calculate the formal charge for each atom in the Lewis structure for H₂SO₄ that we obtained in step (a). For H, O, and S, the formal charges are: \( H: 1 - 1 = 0 \) (Each hydrogen has 0 formal charge.) \( O_{1,2}: 6 - (2 + 6) = -2 \) (The oxygen atoms bonded to hydrogen have -2 formal charge each) \( O_{3,4}: 6 - (2 + 4) = 0 \) (The oxygen atoms not bonded to hydrogen have 0 formal charge.) \( S: 6 - (0 + 4) = +2 \) (Sulfur has a formal charge of +2.)

Step 6: Adjust bonds to minimize formal charge

To minimize the formal charge on the oxygen atoms (O₁ and O₂) bonded to hydrogen, we can convert one lone pair from each of these oxygens into a bond pair with sulfur, creating a double bond (S=O). This will reduce the formal charge on sulfur to zero and on both oxygen atoms to -1.

Step 7: Lewis structure minimizing formal charge

The resulting Lewis structure for H₂SO₄ that minimizes formal charge is: \( H-O-S(O_{3})(O_{4}) \) || | || O_{1} O_{2} Here, O₁ and O₂ have double bonds with sulfur, O₃ and O₄ have single bonds, and all atoms now have formal charges of 0 or -1.

Key Concepts

Sulfuric Acid

Sulfuric acid, with the chemical formula \(\mathrm{H}_{2}\mathrm{SO}_{4}\), is a strong mineral acid widely used in industrial processes. Its structure is arranged with a central sulfur atom surrounded by four oxygen atoms. Two of these oxygen atoms are bonded to hydrogen atoms. This gives sulfuric acid its distinctive tetrahedral geometry. This arrangement is crucial for understanding how the molecule forms and how it behaves in various chemical reactions.

In industrial applications, sulfuric acid is pivotal in producing fertilizers, refining petroleum, and processing metals. Its ability to donate protons makes it an important agent in acid-base reactions.

• Potent industrial chemical
• Central sulfur surrounded by oxygens
• Two oxygen atoms bonded to hydrogen

Understanding the Lewis structure of sulfuric acid helps reveal insights into its reactivity and chemical properties. This structure plays a crucial role in how it interacts with other substances during chemical processes.

Octet Rule

The Octet Rule is a fundamental concept in chemistry that dictates how atoms bond with each other. It states that atoms are more stable when they have eight electrons in their outer shell, mimicking the electron configuration of noble gases.

When drawing the Lewis structure for sulfuric acid to satisfy the octet rule, ensure that the sulfur and oxygen atoms have complete octets. Customarily, the sulfur atom forms bonds with all four oxygen atoms, and each of the oxygen atoms tries to reach an octet by sharing electrons through bonds or keeping lone pairs.

• Atoms stable with 8 electrons in outer shell
• Aims for noble gas electron configuration
• Guides bonding in molecules

While the octet rule helps guide the arrangement of electrons, always remember that some atoms, like sulfur, can have expanded octets due to their ability to hold more electrons in their d-orbitals. This flexibility allows sulfuric acid to maintain its typical structure.

Formal Charge

Formal charge is a concept used to estimate the distribution of electrons in molecules, helping predict the most stable electron arrangement. It is calculated using the formula: \[\text{Formal Charge} = \left(\text{Valence Electrons}\right) - \left(\text{Non-bonding Electrons} + \dfrac{1}{2} \times \text{Bonding Electrons}\right)\]

In sulfuric acid, you calculate the formal charge for each atom to understand its stability. Ideally, structures aim to minimize the formal charge across the molecule.

• Helps predict most stable arrangement
• Aims to minimize charge for stability
• Used for accurate Lewis structures

For sulfuric acid, adjusting bonds to minimize formal charges may involve creating double bonds between sulfur and oxygen. This ensures the molecule has the minimum formal charges possible, leading to a more stable, realistic representation in its Lewis structure.

Valence Electrons

Valence electrons are the outermost electrons in an atom, crucial in determining how an atom will bond with others. For sulfuric acid, the number of valence electrons is derived from each atom forming the molecule. Sulfur, an element in the 16th group of the periodic table, has 6 valence electrons. Each oxygen atom, also in this group, has 6 valence electrons, and each hydrogen atom has 1 valence electron.

Here's how you determine the total valence electrons in \(\mathrm{H}_{2}\mathrm{SO}_{4}\):

• 2 hydrogen atoms: 2 × 1 = 2 valence electrons
• 1 sulfur atom: 6 valence electrons
• 4 oxygen atoms: 4 × 6 = 24 valence electrons

This gives a total of 32 valence electrons to consider when drawing the molecule's Lewis structure, ensuring each atom achieves a complete valence shell wherever possible. Understanding valence electrons is essential for predicting how atoms bond and form stable molecules.