Question

Fluorine reacts with sulfur to form several different covalent compounds. Three of these compounds are \(\mathrm{SF}_{2}, \mathrm{SF}_{4},\) and \(\mathrm{SF}_{6}\) Draw the Lewis structures for these compounds, and predict the molecular structures (including bond angles). Would you expect OF \(_{4}\) to be a stable compound?

Short answer

The Lewis structures of the sulfur-fluorine compounds are as follows: - SF2: : F - S - F : (Linear, bond angle: \(180^{\circ}\)) - SF4: : F - S - F : (Trigonal bipyramidal, axial bond angles: \(90^{\circ}\), equatorial bond angles: \(120^{\circ}\)) || :F: - SF6: Octahedral, bond angles: \(90^{\circ}\) between all F-S-F atoms. OF4 is likely not a stable compound; attempting to form a similar structure to SF4 would result in an expanded octet for the central oxygen atom, violating the octet rule as oxygen does not have accessible d orbitals for bonding.

Step-by-step solution

Lewis structures for each compound

First, we need to draw the Lewis structures for the formulas SF2, SF4, and SF6. We know that sulfur belongs to group 6 (it has 6 valence electrons), and fluorine belongs to group 7 (it has 7 valence electrons). For SF2: - The central atom, sulfur, has 6 valence electrons. - Each fluorine atom contributes 7 valence electrons. On total, the molecule has 20 valence electrons. Form single bonds between sulfur and each fluorine atom. This will use 4 electrons, 16 are left to be placed on the atoms. Fill the octet of each fluorine using the remaining electrons left: : F - S - F : For SF4: - The central atom, sulfur, has 6 valence electrons. - Each fluorine atom contributes 7 valence electrons. On total, the molecule has 34 valence electrons. Form single bonds between sulfur and each fluorine atom. This will use 8 electrons, and 26 are left to be placed on the atom. Fill the octet of each fluorine using the remaining electrons left: : F : || : F - S - F : || :F: For SF6: - The central atom, sulfur, has 6 valence electrons. - Each fluorine atom contributes 7 valence electrons. On total, the molecule has 48 valence electrons. Form single bonds between sulfur and each fluorine atom. This will use 12 electrons, and 36 are left. Fill the octet of each fluorine using the remaining electrons left: : F : | : F-S-F : / \ F F \ / : F - S : \ / F :

Molecular Structures and Bond Angles

Now we will predict the molecular structures and bond angles for each compound: For SF2: - Sulfur is the central atom and has two electron pairs in its outer shell. - The electron pairs form a linear arrangement with an angle of \(180^{\circ}\). - The molecular structure of SF2 is linear with a bond angle of \(180^{\circ}\). For SF4: - Sulfur is the central atom and has four electron pairs in its outer shell. - The electron pairs form a trigonal bipyramidal arrangement. - The fluorine atoms occupy the axial and equatorial positions in the trigonal bipyramidal structure. - The axial bond angles in the SF4 molecule are \(90^{\circ}\) (F-S-F) and the equatorial bond angles are \(120^{\circ}\) (F-S-F). For SF6: - Sulfur is the central atom and has six electron pairs in its outer shell. - The electron pairs form an octahedral arrangement. - The bond angles in SF6 are \(90^{\circ}\) between all of the fluorine atoms (F-S-F).

Stability of OF4

To determine if OF4 is a stable compound, we can compare it to the sulfur-fluorine compounds analyzed above. Oxygen belongs to group 6, similar to sulfur, and has 6 valence electrons. Fluorine belongs to group 7 and has 7 valence electrons. For OF4: - The central atom, oxygen, has 6 valence electrons. - Each fluorine atom contributes 7 valence electrons. On total, the molecule has 34 valence electrons. If we follow the same bonding as in SF4, oxygen will have four single bonds to four fluorine atoms, but also, there will be two remaining lone pairs on the oxygen atom. Thus, oxygen would have an expanded octet exceeding eight electrons—a violation of the octet rule. Whereas sulfur can have an expanded octet due to having available d orbitals, oxygen does not have d orbitals accessible for bonding. This suggests that OF4 would not be a stable compound due to violating the octet rule for oxygen.

Key concepts

Covalent Compounds

Covalent compounds, like the sulfur fluoride examples given (SF₂, SF₄, SF₆), are substances where atoms share electrons through covalent bonds to achieve a fuller outer electron shell—generally regarded as stable with eight electrons. This electron sharing allows each atom in a molecule to attain the electronic configuration of a noble gas, also known as the octet rule.

The number of electrons each atom can share is determined by its valency, or the number of electrons needed to fill or empty its outer shell. For instance, sulfur has a valency of two in SF₂ since it needs two more electrons to complete its octet, whereas in SF₄ and SF₆, the valency increases, allowing for more complex structures due to its ability to house more than eight electrons utilizing d-orbitals.

Molecular Geometry

Molecular geometry is crucial as it affects how molecules interact with each other and react. It is largely determined by the number of electron pairs surrounding the central atom, whether they are bonding pairs or lone pairs. In the exercise, sulfur's geometry changes with each compound due to different electron pair arrangements. SF₂ adopts a linear geometry, while SF₄ shows a more complex trigonal bipyramidal geometry, and SF₆ forms an octahedral arrangement.

Understanding molecular geometry is not just about the position of atoms but also about understanding electron repulsion theory, which tells us that electron pairs will arrange themselves around a central atom as far away from each other as possible to minimize repulsion. This arrangement determines the molecule's shape and bond angles.

Bond Angles

Bond angles provide insight into the three-dimensional arrangement of atoms within a molecule. In simple molecules like SF₂, we have a straightforward bond angle of 180°. However, as molecules become larger and more electrons are involved, bond angles can diversify, as seen in SF₄ and SF₆.

In SF₄, the trigonal bipyramidal structure results in two distinct bond angles—90° and 120° depending on whether the fluorine atoms are in axial or equatorial positions. Meanwhile, SF₆ has a regular octahedral geometry with uniform bond angles of 90°. Differing bond angles can radically affect the molecular properties, such as polarity, which in turn influences intermolecular forces and the compound's physical state.

Octet Rule

The octet rule explains how atoms bond to obtain an outer shell analogous to a noble gas configuration, most often signifying eight electrons. In the textbook solutions, we've seen that sulfur can exceed the octet because it has access to d-orbitals, allowing compounds like SF₆ to form, where sulfur has twelve valence electrons!

However, not all elements can do this—the octet rule is more rigid for lighter elements such as oxygen, which lacks the necessary d-orbitals. This explains why OF₄ is an unlikely molecule. The key takeaway here is that while the octet rule is a fundamental guideline, there are exceptions based on an atom's ability to accommodate electrons beyond the second energy level.