Problem 64 - Chemistry: An Atoms-Focused Approach

Question

Draw Lewis structures that show how electron pairs move and bonds form and break in the following reaction, and identify the Lewis acid and Lewis base. $$ \mathrm{SbF}_{5}(s)+\mathrm{HF}(g) \rightarrow \mathrm{HSbF}_{6}(s) $$

Step-by-Step Solution

Answer

Explain how the bonds form and break in the reaction. Answer: In the reaction between SbF5 and HF, SbF5 acts as the Lewis acid (electron pair acceptor), and HF acts as the Lewis base (electron pair donor). During the reaction, the electron-rich fluorine atom in HF donates an electron pair to the electron-deficient Sb atom in SbF5, resulting in the formation of a new bond between the hydrogen atom of HF and the antimony atom of SbF5. This process leads to the formation of the product HSbF6.

Step 1: Understanding Lewis Acids and Bases

A Lewis acid is a chemical species that can accept an electron pair, whereas a Lewis base is a species that can donate an electron pair. In this reaction, we need to identify which of the reactants acts as a Lewis acid and which acts as a Lewis base.

Step 2: Drawing the Lewis Structures

Firstly, let's draw the Lewis structures of the reactants: SbF5 and HF. For SbF5: Antimony (Sb) has 5 valence electrons, and each of the five Fluorine (F) atoms has 7 valence electrons. In SbF5, the central atom, Sb, is surrounded by five F atoms sharing one electron pair each, forming a single bond. [Sb] <-> [F] || [F] <-> [Sb] <-> [F] || [F] <-> [F] For HF: Hydrogen (H) has 1 valence electron, and Fluorine (F) has 7 valence electrons. They share an electron pair, forming a single bond. [H] <-> [F]

Step 3: Analyzing the Reaction

In the reaction, we have: $$ \mathrm{SbF}_{5}(s)+\mathrm{HF}(g) \rightarrow \mathrm{HSbF}_{6}(s) $$ Notice that, in the product HSbF6, a new bond between Hydrogen (H) and Antimony (Sb) is formed. This reaction can be seen as one of the F atoms in SbF5 being replaced by the H atom from HF. This occurs from the transfer of the electron pair between the atoms involved. HF has a highly polarized bond resulting in the fluorine being electron-rich and able to donate an electron pair. Therefore, it acts as the Lewis base in this reaction. On the other hand, SbF5 can accept electron pairs due to the vacant orbitals, making it a Lewis acid.

Step 4: Writing the Formation and Breaking of Bonds

The reaction can be illustrated via the movement of an electron pair from the F atom of the HF molecule to the SbF5 molecule forming a new bond between the H of the HF and the Sb of the SbF5, giving HSbF6 as the product: [Sb] <-> [F] [H] <-> [F] || . [F] <-> [Sb] /\ .. || - - [Sb] <-> [F] [F] <-> [F] || .. . [Sb] <-> [F] || . [F] <-> [Sb] <-> [F] || [F] <-> [F]

Step 5: Identifying Lewis Acid and Lewis Base

Based on the analysis of the reaction and considering the definition of Lewis acids and bases, we can now conclude that: - SbF5 is the Lewis acid (electron pair acceptor) - HF is the Lewis base (electron pair donor)

Key Concepts

Lewis acid

In the world of chemistry, a Lewis acid is a key concept that refers to any species capable of accepting electron pairs. Imagine it like a magnet ready to attract electrons that are up for grabs.
In the given reaction, we identify SbF$_5$ as the Lewis acid. This is because SbF$_5$ has vacant orbitals which are perfect for accommodating an extra electron pair. These empty orbitals make it seemingly hungry for electrons, driving it to accept them and thereby act as an acid.

• Characteristics of Lewis Acids: They usually have a positive charge or are neutral species with empty orbitals.
• Common Examples: Many transition metals, like Aluminum ions, also serve as Lewis acids.
• Electronegativity: Lower electronegativity often means a stronger Lewis acid, as they are more inclined to gain electrons.

Understanding this concept helps you to easily spot Lewis acids in reactions, sharpening your chemistry skills.

Lewis base

While Lewis acids are on the hunt for electrons, Lewis bases are their exact opposite. They are the electron donors in reactions.
This role is well played by HF in the reaction described. Because HF can provide an electron pair, it qualifies as a Lewis base. Fluorine’s high electronegativity makes it an electron-rich site, promoting its role as a donor.

• Key Traits: Lewis bases usually have lone pairs of electrons, ready to share.
• Common Lewis Bases: Ammonia and hydroxide ions often act as bases in reactions.
• Influence of Electronegativity: Higher electronegativity can enhance the ability to donate electrons, as seen with fluorine in HF.

Recognizing which molecules can act as Lewis bases allows you to predict reactions and understand the underlying chemistry.

electron pairs

In chemical reactions, electron pairs are the heroes moving silently behind the scenes. They are pairs of electrons spinning around and sharing in orbital paths.
In our reaction example, understanding how electron pairs move between HF and SbF$_5$ helps in grasping why bonds form. HF brings an electron pair which it shares with SbF$_5$. This movement is crucial because forming bonds is all about the clever exchange of these pairs.

• Role in Bonding: Electron pairs are the glue that holds atoms together in a molecule.
• Lone Pairs vs. Bonding Pairs: Lone pairs are electron pairs not involved in bonding, while bonding pairs contribute to the formation of chemical bonds.
• Visual Representation: In Lewis structures, electron pairs are often represented by dots or lines between atoms.

By following where electron pairs go, you can predict how new bonds will create compounds and break them down.

bond formation and breaking

Chemistry is often about making and breaking bonds, which is precisely what happens through chemical reactions. Bonding involves the rearrangement of atoms and electron sharing.
In our scenario, the electron pair coming from HF allows for a new bond between H and Sb in SbF$_5$. This results in HSbF$_6$ after the reaction completes. Here, forming bonds leads to a new molecule, highlighting how substances change.

• Formation: Bonds form when electron pairs are shared or transferred between atoms.
• Breaking: Bonds break when these interactions are disrupted, often needing energy input.
• Energy Considerations: Bond formation usually releases energy, while breaking bonds needs energy.

Understanding how bonds are formed and broken unfolds the dynamic nature of molecules, allowing for the vast diversity of chemical reactions we observe.