Problem 100 - Chemistry The Science in Context

Question

Thiosulfuric acid, \(\mathrm{H}_{2} \mathrm{S}_{2} \mathrm{O}_{3},\) can be prepared by the reaction of \(\mathrm{H}_{2} \mathrm{S}\) with \(\mathrm{HSO}_{3} \mathrm{Cl}\) : $$\mathrm{HSO}_{3} \mathrm{Cl}(\ell)+\mathrm{H}_{2} \mathrm{S}(g) \rightarrow \mathrm{HCl}(g)+\mathrm{H}_{2} \mathrm{S}_{2} \mathrm{O}_{3}(\ell)$$ a. Draw a Lewis structure for \(\mathrm{H}_{2} \mathrm{S}_{2} \mathrm{O}_{3},\) given that it is isostructural with \(\mathrm{H}_{2} \mathrm{SO}_{4}\) b. Do you expect \(H_{2} S_{2} O_{3}\) to be a stronger or weaker acid than \(\mathrm{H}_{2} \mathrm{SO}_{4} ?\) Explain your answer.

Step-by-Step Solution

Answer

Answer: Thiosulfuric acid (H2S2O3) is a weaker acid than sulfuric acid (H2SO4) because the conjugate base of sulfuric acid has more electronegative atoms (oxygen) stabilizing the negative charge, making it a more stable conjugate base and thus a stronger acid.

Step 1: Determine the total number of valence electrons

First, we need to determine the number of valence electrons for each atom present in \(H_{2}S_{2}O_{3}\). Hydrogen has 1 valence electron, sulfur has 6, and oxygen has 6. So, in total, the molecule has \((2 \times 1) + (2 \times 6) + (3 \times 6) = 28\) valence electrons.

Step 2: Arrange the atoms in the structure

Since the molecule is isostructural with \(H_{2}SO_{4}\), we can use the same arrangement with sulfur atoms in the center, oxygen atoms surrounding the sulfur atoms, and hydrogen atoms bonded to the outer oxygen atoms. The structure will look like this: \(-O-S-S-O-\), with hydrogens bonded to the outer oxygen atoms.

Step 3: Connect atoms with single bonds

Next, we need to connect the atoms with single bonds and distribute the remaining valence electrons. Each of the single bonds uses 2 electrons, so we will have used \((2 \times 3) + (2 \times 1) = 10\) electrons so far. There are 18 electrons left to distribute.

Step 4: Complete the octet structure for each atom

Using the remaining 18 electrons, we will complete the octet structure for each atom. The oxygens and sulfurs should have 8 electrons each, while hydrogen should have 2 electrons each. After distributing the remaining electrons, each oxygen and sulfur atom achieves an octet structure, and each hydrogen atom has 2 electrons.

Step 5: Verify the Lewis structure

Finally, we need to verify that our Lewis structure for \(H_{2}S_{2}O_{3}\) is correct: - Hydrogen: 1 bond and 0 lone pairs - Sulfur: 2 bonds and 2 lone pairs - Oxygen: 2 bonds and 2 lone pairs Since all atoms have achieved their respective desired electron configurations, the Lewis structure for \(H_{2}S_{2}O_{3}\) is correct. #b. Acid Strength of H2S2O3 Compared to H2SO4#

Step 1: Identify the acidic protons

First, we should identify which protons are acidic in both compounds. In \(H_{2}S_{2}O_{3}\) and \(H_{2}SO_{4}\), the acidic protons are those bonded to the oxygen atoms.

Step 2: Analyze the stability of the conjugate base

A good way to compare acid strengths is to analyze the stability of their conjugate bases. A more stable conjugate base is formed by a stronger acid. When an acidic proton is removed from \(H_{2}S_{2}O_{3}\), the conjugate base formed has negatively charged sulfur atoms, while the conjugate base of \(H_{2}SO_{4}\) has negatively charged oxygen atoms.

Step 3: Compare electronegativities of atoms in conjugate bases

Now, we need to compare the electronegativities of sulfur and oxygen. Oxygen is more electronegative than sulfur, which means it is better at stabilizing the negative charge in the conjugate base.

Step 4: Determine the stronger acid

Since the conjugate base of \(H_{2}SO_{4}\) has more electronegative atoms stabilizing the negative charge, it is more stable and thus \(H_{2}SO_{4}\) is a stronger acid compared to \(H_{2}S_{2}O_{3}\). In conclusion, \(H_{2}S_{2}O_{3}\) is a weaker acid than \(H_{2}SO_{4}\) due to the difference in electronegativity between sulfur and oxygen atoms in their respective conjugate bases.

Key Concepts

Lewis structure

A Lewis structure is like a map of a molecule showing how the atoms are arranged and connected. To draw the Lewis structure for thiosulfuric acid, \(H_{2}S_{2}O_{3}\), start by counting the valence electrons. Each hydrogen (H) has 1 valence electron, while both sulfur (S) and oxygen (O) have 6. In total, we have \(2 \times 1 + 2 \times 6 + 3 \times 6 = 28\) valence electrons.

Picture sulfur atoms in the center with oxygen atoms surrounding them. Hydrogen atoms connect to the outer oxygen atoms. The key is to connect the atoms using single bonds first, which consumes electrons. Any leftover electrons complete the "octet" — which means providing 8 valence electrons — for each of the S and O atoms.

Checking ensures every atom satisfies its ideal electron configuration, making our Lewis structure accurate.

conjugate base stability

The stability of a conjugate base gives insight into the strength of its acid. When an acid loses a proton \((H^+)\), it forms what we call a conjugate base. A stable conjugate base points to a stronger acid, as it efficiently handles the negative charge from losing a proton.

For thiosulfuric acid \(H_{2}S_{2}O_{3}\) and sulfuric acid \(H_{2}SO_{4}\), we examine their conjugate bases. The removal of a proton from thiosulfuric acid leaves sulfur atoms bearing negative charge, whereas from sulfuric acid, oxygen atoms bear this negative charge. Because oxygen is better at stabilizing negative charges than sulfur, sulfuric acid forms a more stable conjugate base, proving it is a stronger acid.

Understanding which atom takes the negative charge affects acid strength.

valence electrons

Valence electrons are the outermost electrons of an atom. They dictate how atoms bond together to create molecules. When analyzing \(H_{2}S_{2}O_{3}\), it's key to know how many valence electrons each atom possesses, allowing us to sketch the Lewis structure accurately.

Hydrogen has 1, sulfur 6, and oxygen 6 valence electrons. In total, \(H_{2}S_{2}O_{3}\) includes 28 valence electrons. These electrons are used in bonds between atoms and as lone pairs — electrons not involved in bonding — to satisfy the "octet rule," where atoms aim for 8 electrons in their outer shell.

Using them efficiently aids in creating stable structures and understanding chemical reactions.

electronegativity

Electronegativity is an atom's tendency to attract electrons in a bond. It's a crucial idea when comparing the strength of different acids.

In the context of \(H_{2}S_{2}O_{3}\) and \(H_{2}SO_{4}\), this concept crucially affects how we compare acid strengths. Oxygen, being more electronegative than sulfur, is better at drawing and holding onto negative charge when a molecule loses a hydrogen ion. So, \(H_{2}SO_{4}\) has a stronger acidic character due to the ability of its conjugate base (formed when an acidic proton is removed) to stabilize the charge effectively.

Electronegativity, therefore, often dictates which bonds form and how molecules behave in reactions.

sulfur and oxygen atoms

Understanding the difference between sulfur and oxygen atoms is key when studying molecules like \(H_{2}S_{2}O_{3}\) and \(H_{2}SO_{4}\). These two elements share some chemical properties but differ significantly in their electronegativity.

Oxygen is more electronegative than sulfur, meaning it can better stabilize the negative charge in chemical reactions. This results in sulfuric acid being stronger since its oxygen-containing conjugate base can more efficiently bear the change after an acidic proton is lost.

The positioning of sulfur and oxygen in the periodic table explains these differences, with sulfur sitting directly below oxygen. Knowing these differences helps predict molecular behavior and reactivity, critical in fields like organic chemistry.