Draw the Lewis structure for the SnF62- ion.

Draw the Lewis structure for the SnF62- ion.

The Correct Answer and Explanation is:

Lewis Structure for the SnF₆²⁻ Ion:

Step-by-step construction:

1. Count total valence electrons:
   • Tin (Sn) is in Group 14 → 4 valence electrons
   • Each fluorine (F) atom is in Group 17 → 7 valence electrons × 6 = 42
   • The ion has a 2⁻ charge → add 2 more electrons
   • Total = 4 + 42 + 2 = 48 valence electrons
2. Choose the central atom:
   • Sn is less electronegative than F, so Sn is the central atom
3. Place 6 fluorine atoms around the central Sn atom:

rCopyEdit        F
        |
    F - Sn - F
        |
        F

• All 6 fluorine atoms are bonded to Sn

4. Place a single bond (2 electrons) between Sn and each F:
   • 6 bonds × 2 electrons = 12 electrons used
   • Remaining electrons = 48 – 12 = 36 electrons
5. Complete the octets for each fluorine:
   • Each fluorine needs 6 more electrons to complete its octet
   • 6 fluorine atoms × 6 electrons = 36 electrons
   • All 36 electrons used, none left
6. Check formal charges:
   • Sn forms 6 bonds → it uses all 4 of its valence electrons and borrows more from bonding
   • Fluorine atoms each have 3 lone pairs and 1 bond → 7 valence electrons accounted for
   • The entire molecule has a 2⁻ charge which matches the ion
7. Add brackets and charge to show it’s an ion:

rCopyEdit       ..
   :F:      :F:
    |        |
:F:–Sn–F:  
    |        |
   :F:      :F:
       ..
[SnF₆]²⁻

Explanation

The Lewis structure of the hexafluorostannate ion, SnF₆²⁻, illustrates how atoms share electrons to satisfy the octet rule while accounting for the overall charge. To construct it, we begin by calculating the total number of valence electrons. Tin contributes 4, and each of the six fluorine atoms contributes 7, resulting in 46 electrons. Since the ion carries a 2⁻ charge, we add 2 more, giving us 48 valence electrons in total.

Next, we identify the central atom. Tin (Sn) is less electronegative than fluorine and has the ability to expand its valence shell due to its position in Period 5. Thus, it becomes the central atom with six fluorine atoms arranged symmetrically around it.

We then form single bonds between Sn and each F. These 6 single bonds consume 12 electrons. The remaining 36 electrons are used to complete the octets of the surrounding fluorine atoms by placing three lone pairs (6 electrons) on each.

After placing all 48 electrons, we check the octets. Each fluorine has 8 electrons: 6 from lone pairs and 2 from the Sn–F bond. Tin forms 6 bonds, and while it exceeds the octet, this is acceptable for elements in Period 5 or beyond due to available d orbitals.

Finally, we enclose the structure in brackets to denote it is an ion and indicate the 2⁻ charge in the upper right corner. This complete Lewis structure helps us understand the molecule’s bonding, electron distribution, and overall charge, which is important for predicting reactivity and molecular geometry.