Question

Choose the stronger acid: (a) \(\mathrm{HBr}\) or \(\mathrm{HCl}\), (b) \(\mathrm{H}_{2} \mathrm{O}\) or \(\mathrm{HF},\) (c) \(\mathrm{H}_{2} \mathrm{~S}\) or \(\mathrm{HBr}\). Give your reasons.

Short answer

The stronger acids are: (a) HBr, (b) HF, (c) HBr. This is based on atom size for halogens (larger halogens form stronger acids in hydrohalic acids) and electronegativity (more electronegative atoms form stronger acids).

Step-by-step solution

Compare Acid Strength Based on Halogens

Acid strength increases with the size of the halogen atom for hydrohalic acids. Compare the size of bromine (Br) and chlorine (Cl). Bromine is larger than chlorine, which means that HBr has a weaker H-X bond than HCl and can release a proton more easily.

Compare Acidity of Water and Fluoride Acid

The strength of an acid can be affected by the electronegativity of the atom to which the hydrogen is bonded. Fluorine is highly electronegative, making HF a stronger acid than H2O because the H-F bond is more polar and more prone to donate a proton.

Compare Acidity between H2S and HBr

H2S is a binary acid where the hydrogen is bonded to sulfur, whereas HBr is a hydrohalic acid. Since bromine is a halogen and sulfur is a non-metal, generally hydrohalic acids like HBr will be stronger acids than binary acids like H2S.

Key concepts

Hydrohalic Acids

Hydrohalic acids are a subset of acids composed of hydrogen and halogen atoms. The most common hydrohalic acids are hydrofluoric acid (HF), hydrochloric acid (HCl), hydrobromic acid (HBr), and hydroiodic acid (HI). A fundamental aspect of understanding hydrohalic acids involves assessing their strength as acids.

One key factor in determining the strength of hydrohalic acids is the size of the halogen atom involved. As we progress down the halogen group in the periodic table, the atoms increase in size – fluorine being the smallest and iodine the largest. This size increment leads to a looser hold on the bond between the halogen and hydrogen atoms. Consequently, larger halogens form weaker bonds with hydrogen, rendering acids like HBr and HI stronger – they can more readily donate a hydrogen proton.

In the given exercise, we compared HBr and HCl. Bromine is a larger atom than chlorine, thus HBr has a weaker H-Br bond compared to the H-Cl bond in HCl, making it easier for HBr to donate a proton and thereby act as a stronger acid.

Acid Electronegativity

Electronegativity is a chemical property that describes the tendency of an atom to attract electrons towards itself in a chemical bond. When it comes to acids, the electronegativity of the atoms to which hydrogen atoms are bonded can significantly influence acid strength.

In the case of hydrofluoric acid (HF), we observe a high electronegativity due to the fluorine atom, which heavily attracts the bonding electrons. This electronegative pull increases the polarity of the H-F bond, which in turn enhances the ability of the acid to donate a proton in aqueous solutions. This concept explains the stronger acidity of HF compared to water (H2O) as the oxygen atom is less electronegative than fluorine.

When examining H2O and HF, it’s clear that the more polar bond in HF (due to fluorine's high electronegativity) leads to easier proton donation, resulting in a more robust acid. This level of understanding allows us to predict the relative acid strengths when comparing molecules with different electronegativities attached to the hydrogen atom.

Proton Donation

The ability of an acid to donate a proton, which is a hydrogen atom without its electron, is a defining feature of its strength. Proton donation is fundamental to the concept of acidity in chemistry, and it is the basis upon which the Bronsted-Lowry definition of acids and bases is founded.

An acid is considered strong if it readily donates protons in an aqueous solution, leading to a high concentration of hydronium ions (H3O+). This willingness to give up a proton is influenced by several factors, including bond strength, the atom's electronegativity to which the hydrogen is attached, and the solvent in which the acid is dissolved.

In comparing H2S and HBr, it is the weaker bond due to the larger size of bromine, along with its reasonable electronegativity, that makes HBr the stronger acid compared to H2S. HBr donates protons more easily than H2S due to the weaker H-Br bond, highlighting the central role that proton donation plays in evaluating acid strength.