Question

The decreasing order of power of boron halides to act as Lewis acids is (a) \(\mathrm{BF}_{3}>\mathrm{BCl}_{3}>\mathrm{BBr}_{3}\) (b) \(\mathrm{BBr}_{3}>\mathrm{BCl}_{3}>\mathrm{BF}_{3}\) (c) \(\mathrm{BCl}_{3}>\mathrm{BF}_{3}>\mathrm{BBr}_{3}\) (d) \(\mathrm{BCl}_{3}>\mathrm{BBr}_{3}>\mathrm{BF}_{3}\)

Short answer

The correct order of decreasing Lewis acidity for boron halides is BBr3 > BCl3 > BF3.

Step-by-step solution

Understand the Concept of Lewis Acidity

Lewis acids are species that can accept an electron pair. Boron halides (BX3, where X is a halogen) act as Lewis acids by accepting lone pairs from Lewis bases. The Lewis acidity of boron halides depends on the ability of the boron atom to accept electron pairs, which is influenced by the electronegativity and size of the halogen atoms attached to it.

Identify the Influence of Electronegativity

The electronegativity of the halogens decreases as we move down the group in the periodic table (F > Cl > Br). A more electronegative atom will draw electron density towards itself and away from the boron atom, making the boron less electrophilic and thus less able to act as a Lewis acid.

Identify the Influence of Halogen Atom Size

As we move down the group in the periodic table, the size of the halogens increases (F < Cl < Br). Larger halogen atoms are less effective at overlapping their orbitals with those of the boron, leading to less effective back-donation of electron density and hence increased Lewis acidity of the boron center.

Combine the Effects of Electronegativity and Size to Determine Order

The smaller and more electronegative the halogen, the more it withdraws electron density from the boron atom, decreasing its Lewis acidity. Thus, as fluorine is the most electronegative and smallest, BF3 will be the least Lewis acidic, while BBr3, with the least electronegative and largest halogen, will be the most Lewis acidic. Therefore, the order of Lewis acidity is BBr3 > BCl3 > BF3.

Key concepts

Boron Halides

Boron halides are a group of compounds that consist of a boron atom bonded to three halogen atoms, represented as BX₃, where X can be a halide such as fluorine (F), chlorine (Cl), or bromine (Br). These compounds are known to act as Lewis acids because the boron atom has an empty p-orbital which can accept an electron pair from a Lewis base.

The ability of boron halides to act as Lewis acids makes them incredibly valuable in numerous chemical reactions, especially in organic synthesis where they are used as catalysts to accelerate reactions. By understanding the properties of these substances, chemists can predict the reactivity and interaction with various Lewis bases, which is critical in designing effective synthetic routes in the laboratory.

Electronegativity

Electronegativity is a fundamental concept in chemistry that describes the tendency of an atom to attract electrons to itself when it's part of a molecule. On the Pauling scale, fluorine (F) is the most electronegative element, with a value of around 4.0, and the values generally decrease as you move down a group in the periodic table.

When dealing with boron halides, the differing electronegativities of the halogens affect the electron density around the boron atom. A higher electronegative halogen, like fluorine in BF₃, will pull more electron density towards itself, reducing the ability of the boron atom to accept electrons and act as a Lewis acid. This detail is key to predicting the reactivity of boron halides in chemical processes.

Periodic Table

The periodic table is an essential tool for chemists, organizing elements by increasing atomic number and grouping them by similar properties. The periodic trends, such as electronegativity and atomic size, play a crucial role in understanding the behavior of elements.

For example, as you move down Group 13, where boron is located, the elements have an increasing tendency to accept electron pairs due to the growing atomic size; this influences their Lewis acidity. The trends seen in Group 17 (the halogens) are also noteworthy, as these trends will affect the characteristics of boron halides.

Lewis Acidity Order

The Lewis acidity order for boron halides can be determined by considering both electronegativity and atomic size of the attached halogens. Boron trihalides, such as BF₃, BCl₃, and BBr₃, vary in their ability to accept electron pairs based on the nature of the halogen substituent.

Considering the effects of electronegativity and halogen size, we find that BBr₃ is the most potent Lewis acid in the group, as bromine has the lowest electronegativity and the largest size, making it less able to attract the electron density around boron. Conversely, BF₃ is the least potent because fluorine's high electronegativity pulls electron density towards itself, making the boron less electrophilic.

Electronic Configurations

Understanding electronic configurations of atoms allows chemists to predict and explain chemical reactivity and bonding patterns. Boron, for example, has an electronic configuration of 1s² 2s² 2p¹, which leaves it with an empty p-orbital that can act as an electron pair acceptor or Lewis acid.

The halogens have varying electronic configurations, but all have seven valence electrons, which they can share with boron's empty p-orbital. The efficiency of this overlap, and hence the Lewis acidity, is dependent not just on the halogen's electronegativity, but also on the orbital size and the capacity for effective orbital interactions with the vacant boron p-orbital.