Question

Which of the following are Lewis acids? (1) \(\mathrm{BF}_{3}\) (2) \(\mathrm{H}_{2} \mathrm{O}\) (3) \(\mathrm{HSO}_{4}^{-}\) (4) \(\mathrm{SO}_{3}\) (a) (1) and (3) (b) (1) and (2) (c) (1) and (4) (d) (3) and (4)

Short answer

The Lewis acids are \(\mathrm{BF}_3\) and \(\mathrm{SO}_3\); choose (c).

Step-by-step solution

Understanding Lewis Acids

A Lewis acid is defined as a chemical species that can accept an electron pair. To identify Lewis acids, we look for molecules or ions that have an electron-deficient atom that can accept electrons.

Evaluating BF3

\(\mathrm{BF}_3\) is a boron trifluoride molecule with boron as the central atom. Boron has only six electrons in its valence shell, making it electron-deficient and able to accept a pair of electrons. Therefore, \(\mathrm{BF}_3\) is a Lewis acid.

Evaluating H2O

Water \(\mathrm{H}_2\mathrm{O}\) has a complete octet and its central atom, oxygen, has two lone pairs of electrons. While \(\mathrm{H}_2\mathrm{O}\) can donate electron pairs and act as a Lewis base, it does not act as a Lewis acid since it cannot accept electron pairs readily compared to known Lewis acids.

Evaluating HSO4-

The bisulfate ion \(\mathrm{HSO}_4^-\) is not electron deficient. It usually acts as a base since it has the capacity to donate electrons (for instance, donating a proton to become \(\mathrm{SO}_4^{2-}\)). It does not typically act as a Lewis acid by accepting an electron pair.

Evaluating SO3

\(\mathrm{SO}_3\) is sulfur trioxide, and the sulfur atom can form additional bonds by accepting an electron pair. The sulfur atom in \(\mathrm{SO}_3\) has empty d orbitals, which makes it capable of accepting electrons, classifying it as a Lewis acid.

Comparing Options

Based on the evaluation: \(\mathrm{BF}_3\) and \(\mathrm{SO}_3\) are Lewis acids. None of the other listed substances act as Lewis acids. Thus, correct options include molecules (1) and (4), corresponding to answer (c).

Key concepts

Electron Pair Acceptance

One fundamental aspect of the Lewis acid-base theory is the concept of electron pair acceptance. In a chemical reaction, a Lewis acid is any entity that can accept an electron pair. This characteristic stems from the presence of an electron-deficient atom within the molecule or ion. For instance, consider boron trifluoride \(\mathrm{BF}_3\). Boron, the central atom, has only six electrons in its outer shell. Hence, it can accept an electron pair from another species, filling its valence shell.

It’s crucial to understand this electron deficiency because it directly influences a substance's ability to participate in chemical reactions as a Lewis acid. Not all atoms or molecules can accept electrons easily; they need to be electron-deficient to do so.

• Electron-deficient atoms have fewer electrons than required for a stable configuration.
• These atoms seek electrons to achieve a filled valence shell or satisfy the octet rule.

Thus, identifying whether a compound can act as a Lewis acid largely depends on an evaluation of its electron pair acceptance potential.

Lewis Acid-Base Theory

Lewis acid-base theory broadens the traditional concept of acids and bases beyond just hydrogen or hydroxide ion involvement. In this theory:

• Lewis acids are substances that can accept an electron pair.
• Lewis bases are substances that donate an electron pair.

This theory covers a wider range of reactions than the typical Brønsted-Lowry or Arrhenius definitions. While the latter definitions rely on the presence of hydrogen or hydroxide ions, Lewis theory focuses on electron exchanges.

By focusing on electrons, Lewis acid-base theory explains complex reactions, especially those involving coordination compounds and reactions where no protons are transferred. For instance, in evaluating \(\mathrm{SO}_3\), sulfur trioxide can accept an electron pair due to its empty d orbitals, making it a suitable Lewis acid.

Chemical Species Evaluation

Evaluating a chemical species to determine its role as a Lewis acid involves analyzing its electron structure and the presence of electron-deficient atoms. This process includes considering:

• The valence electrons around critical atoms.
• Electronegativity and the resulting polarity of bonds.
• The presence of empty or half-filled orbitals.

Taking \(\mathrm{BF}_3\) (boron trifluoride) as an example, the boron's electron configuration demonstrates a clear deficiency, making it an ideal candidate for Lewis acidity. In contrast, water \(\mathrm{H}_2\mathrm{O}\), with a full octet around oxygen, does not have the capacity for electron pair acceptance and thus does not function as a Lewis acid.

The process also involves understanding molecular geometry. In sulfur trioxide \(\mathrm{SO}_3\), the planar trigonal structure around the sulfur atom facilitates interactions where electron donation from another molecule completes its electron arrangement.

Molecular Geometry and Electrons

Molecular geometry plays a critical role in determining whether a chemical species can act as a Lewis acid. The shape of a molecule affects the distribution of electrons and the availability of electron-deficient sites. Let’s look at some characteristics:

• Planar structures often allow greater accessibility to electron-deficient atoms.
• Molecules with vacant orbitals like d orbitals provide spaces for new bonds.

For example, in \(\mathrm{SO}_3\) (sulfur trioxide), the central sulfur atom is surrounded by oxygen atoms in a trigonal planar arrangement. Its capacity to form additional covalent bonds by accepting electron pairs hinges on the geometry, making it an effective Lewis acid.

Contrastingly, molecular geometry in \(\mathrm{H}_2\mathrm{O}\) (water), with its bent shape and complete octet around oxygen, does not facilitate additional electron pair acceptance. Therefore, understanding the spatial arrangement of atoms helps predict the chemical behavior and reactivity applicable to Lewis acids and bases.