Question

Write Lewis structures for the following species, both of which involve coordinate covalent bonding: (a) tetrafluoroborate ion, \(\mathrm{BF}_{4}^{-}\), used in metal cleaning and in electroplating baths. (b) boron trifluoride ethylamine, used in curing epoxy resins (ethylamine is \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NH}_{2}\) ).

Short answer

The tetrafluoroborate ion has a Lewis structure with the boron connected to four fluorine atoms with single bonds, one lone pair on B, and three lone pairs on each F. Boron trifluoride ethylamine's Lewis structure involves drawing \(\mathrm{BF}_{3}\) and ethylamine separately, then showing a coordinate bond from N to B.

Step-by-step solution

Tetrafluoroborate ion \(\mathrm{BF}_{4}^{-}\) Lewis structure

First, draw the atoms and connect them with a single bond. Boron (B) is in the center with a fluorine (F) atom on each side, a total of four. Now, count the valence electrons. B has 3 and each four F atom has 7, adding to 31. To satisfy the octet rule, each F gets three pairs of electrons, so we have a total of \(4x6=24\) electrons in these electron pairs around F atoms. The remaining \(31-24=7\) electrons are placed as lone pairs around the central B atom. One extra electron from the negative charge is also added to the B atom. Thus, there is one lone pair on B and three lone pairs on each F.

Boron trifluoride ethylamine Lewis structure

You must create the Lewis structure for both boron trifluoride (\(\mathrm{BF}_{3}\)) and ethylamine (\(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NH}_{2}\)) separately first. \(\mathrm{BF}_{3}\) has 24 valence electrons, each B-F bond uses 2 electrons and the remaining 6 are placed on F atoms. Only three valence electrons are used by B in \(\mathrm{BF}_{3}\) so it is electron-deficient and able to accept a pair from another molecule. Ethylamine is very similar to ammonia, with an NH2 attached to a C2H5; it has one lone pair on the N that can be donated. Draw each molecule separately, then combine them, showing a coordinate bond from N to B.

Understanding the coordinate covalent bond

Reiterate that in both compounds, there is a coordinate covalent bond where a lone pair of electrons from one atom is donated to a electron-deficient atom to form a bond. This forms the compound structure.

Key concepts

Coordinate Covalent Bonding

Understanding coordinate covalent bonding is crucial when dealing with complex molecules. This type of bonding occurs when one atom donates a pair of electrons to another atom to form a bond. Unlike a regular covalent bond, where each atom supplies one electron, in a coordinate bond, both electrons come from the same atom.

For example, in the case of boron trifluoride ethylamine, the nitrogen (N) atom of ethylamine has a lone pair of electrons. Boron trifluoride (BF₃) is electron-deficient and does not adhere to the octet rule, making it an ideal electron pair acceptor. Thus, the nitrogen donates its two electrons to the boron, forming a coordinate covalent bond.

Tetrafluoroborate Ion

The tetrafluoroborate ion, BF₄⁻, is an example of a molecule where symmetry and coordinate covalent bonding come into play. Each fluorine atom donates a pair of electrons to the central boron atom. Despite the negative charge indicating an additional electron, boron in BF₄⁻ still does not have a full octet. The negative charge is delocalized, meaning it's spread out across the entire ion rather than residing on a single atom.

When drawing the Lewis structure, it's essential to represent the formal charges correctly. All four fluorine atoms carry a formal charge of zero, while the boron atom has a formal charge of -1 due to the extra electron associated with the negative ion.

Boron Trifluoride Ethylamine

Boron trifluoride ethylamine is a product of a coordinate covalent bond between a boron trifluoride molecule and an ethylamine. The Lewis structure for each molecule needs to be drawn independently before they can be combined. Boron trifluoride is a planar molecule with three fluorine atoms bonded to a central boron atom. This molecule is incomplete in terms of the octet rule, which allows it to accept a pair of electrons from a donor like ethylamine.

When ethylamine combines with BF₃, the nitrogen's lone pair fills the electron deficiency in boron. This results in a stable compound used in applications such as curing epoxy resins.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in chemical bond formation. They determine the reactivity and bonding capability of the atom. In the given examples, the valence electrons of boron and fluorine play a pivotal role. Boron, with only three valence electrons, does not satisfy the octet rule natively but can form compounds through coordinate bonding.

In the case of the tetrafluoroborate ion, counting the valence electrons correctly and understanding their distribution is key to drawing an accurate Lewis structure. The presence of valence electrons on surrounding fluorine atoms and the central boron provides insight into the molecule's bonding and reactivity.

Octet Rule

The octet rule is a chemical rule of thumb that reflects the observation that atoms of main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. While this rule is generally followed, there are exceptions. Boron trifluoride, for example, has only six valence electrons around the boron atom and violates the octet rule.

It is this violation that allows for the formation of coordinate covalent bonds as seen with the boron trifluoride ethylamine complex. The octet rule guides us in predicting the bonding patterns in most molecules but it's equally important to recognize its limitations and the existence of compounds that deviate from this rule.