Question

Explain the following observations: (a) \(\mathrm{HCl}\) is a stronger acid than \(\mathrm{H}_{2} \mathrm{~S}\); (b) \(\mathrm{H}_{3} \mathrm{PO}_{4}\) is a stronger acid than \(\mathrm{H}_{3} \mathrm{AsO}_{4}\); (c) \(\mathrm{HBrO}_{3}\) is a stronger acid than \(\mathrm{HBrO}_{2}\); (d) \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) is a stronger acid than \(\mathrm{HC}_{2} \mathrm{O}_{4}^{-}\); (e) benzoic acid \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\right)\) is a stronger acid than phenol \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{OH}\right)\).

Short answer

(a) HCl is a stronger acid than H2S because Cl is more electronegative than S, making it easier for the H+ to be donated. (b) H3PO4 is a stronger acid than H3AsO4 because P is more electronegative than As, allowing for easier proton donation. (c) HBrO3 is a stronger acid than HBrO2 due to the extra oxygen atom, which increases the stability of the conjugate base. (d) H2C2O4 is a stronger acid than HC2O4- because the first proton loss results in a more stable conjugate base due to resonance. (e) Benzoic acid is a stronger acid than phenol because the conjugate base of benzoic acid is more stabilized by resonance.

Step-by-step solution

(a) Compare the acidity of HCl and H2S#acidity ofolutinnee

The acidity of an acid depends on its ability to donate a proton (H+) to a solution. In this case, we should compare the strength of the hydrogen bond in both molecules. HCl is a stronger acid than H2S because Cl is more electronegative than S, meaning that it attracts the bonding electrons more, weakening the H-Cl bond and making it easier for the H+ to be donated to a solution. Additionally, the size difference between H2S (larger) and HCl (smaller) also contributes to the higher electronegativity/enabling HCl to be a stronger acid.

(b) Compare the acidity of H3PO4 and H3AsO4

Again, we must compare the bonds between the hydrogen atoms and the central atom (P or As). Phosphorus (P) is more electronegative than Arsenic (As), resulting in a stronger attraction to the hydrogen atoms, and thus, a weaker bond allowing for the hydrogen to be donated more easily. H3PO4 is, therefore, a stronger acid than H3AsO4.

(c) Compare the acidity of HBrO3 and HBrO2

The acidity of oxyacids like HBrO3 and HBrO2 depends on the number of oxygen atoms attached to the central atom, which influences the stability of the conjugate base. HBrO3 has more oxygen atoms (3) compared to HBrO2 (2). The extra oxygen atom increases the stability of the conjugate base, making HBrO3 a stronger acid than HBrO2 because it more easily loses a proton.

(d) Compare the acidity of H2C2O4 and HC2O4-

H2C2O4 is an oxalic acid with two acidic protons. When one proton is lost, it forms HC2O4-. At this point, the negative charge on the conjugate base can be delocalized across oxygen atoms, stabilizing it. H2C2O4 is a stronger acid than HC2O4- because the first proton loss results in a conjugate base that is more stable due to resonance, while the second proton loss results in a doubly charged anion, which is less stable due to electrostatic repulsion.

(e) Compare the acidity of benzoic acid and phenol

Both benzoic acid and phenol have a carboxyl group, but benzoic acid has an additional carboxyl group (COOH), which is the acidic site. The resonance effect stabilizes the conjugate base of benzoic acid, as the negative charge is delocalized across the benzene ring. Phenol is less acidic because the negative charge on the conjugate base is not as delocalized and stabilized by resonance compared to benzoic acid. Therefore, benzoic acid is a stronger acid than phenol.

Key concepts

Electronegativity

Electronegativity refers to the ability of an atom to attract electrons towards itself. This concept is crucial in understanding acid strength, particularly when comparing acids like HCl and H2S. When chlorine (Cl) and sulfur (S) bond with hydrogen (H), the differences in electronegativity lead to differences in acid strength.
Chlorine is more electronegative than sulfur, which means it attracts the electrons in the H-Cl bond more strongly than sulfur does in the H-S bond. This stronger attraction weakens the H-Cl bond, making it easier for the hydrogen ion (H+) to dissociate, thereby making HCl a stronger acid than H2S.

• The greater the electronegativity, the weaker the bond with hydrogen, leading to more potent acid behavior.
• Similarly, substances like H3PO4 are stronger acids than H3AsO4 due to phosphorus being more electronegative than arsenic, allowing for greater H+ donation.

Oxyacids

Oxyacids are acids that contain hydrogen, oxygen, and another element (the central atom). The acidity of an oxyacid is often influenced by the number of oxygen atoms attached to the central atom. This aspect is evident when comparing acidity between HBrO3 and HBrO2.
In general, the more oxygen atoms present, the stronger the acid. This is because additional oxygen atoms increase the overall electronegativity, which helps to stabilize the conjugate base after the acid donates a proton. Thus, HBrO3, with three oxygen atoms, is a stronger acid than HBrO2, which has two.

• More oxygen atoms mean greater stability of the conjugate base due to the enhancement of resonance.
• The effect of additional oxygen atoms is comparable to the tug-of-war effect, where added players (oxygen) enhance the central atom's pull on electrons.

Resonance Stability

Resonance stability relates to how well a conjugate base can distribute its charge over multiple atoms. This is a key factor determining the strength of acids such as benzoic acid and oxalic acid (H2C2O4).
When an acid donates a proton, it converts into its conjugate base. If this conjugate base can spread out its negative charge through resonance, it becomes more stable and thus strengthens the acidity of the original acid. For example, benzoic acid forms a conjugate base that resonates over the benzene ring, making it more stable than phenol's conjugate base.

• Resonance in conjugate bases leads to greater delocalization of electrons, enhancing stability.
• Oxalic acid further demonstrates this; its first proton loss leads to a conjugate base with resonating negativity, enhancing acid strength.

Conjugate Base

A conjugate base forms when an acid donates a proton. The stability of this conjugate base directly affects the acid's strength. A more stable conjugate base allows for easier proton donation, thus a stronger acid.
For instance, when HBrO3 donates a proton, the resulting conjugate base is more stable due to the presence of extra oxygen atoms that stabilize the negative charge through resonance. Similarly, when benzoic acid loses a proton, its conjugate base stabilizes the negative charge across the benzene ring, making it a stronger acid than phenol.

• Stable conjugate bases favor the forward reaction of proton loss, strengthening acid characteristics.
• Increased stabilization through resonance or electronegative atoms directly contributes to a stronger conjugate base.

Proton Donation

Proton donation is the process through which an acid releases a hydrogen ion (H+). This is the hallmark of acidic behavior. Strong acids readily donate protons, which is influenced by factors like bond strength and conjugate base stability.
HCl is a prime example due to its shorter and more polarized bond compared to H2S. This makes H+ donation easier for HCl, making it a more potent acid. Similarly, substances that can stabilize their conjugate base through resonance or additional electronegative atoms also enhance their capacity for proton donation.

• Acid strength increases with ease of proton donation; weaker bonds and stable conjugate bases facilitate this.
• Efficient resonance or electronegative balancing enhances ability for protons to dissociate.