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In the \(\mathrm{NH}_{4}+\) ion, there is one nitrogen atom and four hydrogen atoms. Nitrogen has 5 valence electrons, and hydrogen has 1 valence electron. However, since the ion has a charge of +1, we need to subtract one electron from the total count. So, the total number of valence electrons available is: Total valence electrons = (1 nitrogen x 5) + (4 hydrogen x 1) - 1 (charge) = 5 + 4 - 1 = 8 electrons.

Draw a single bond between the nitrogen atom and each hydrogen atom. These 4 single bonds utilize 8 electrons (2 electrons per single bond).

There are no remaining valence electrons after connecting the central nitrogen atom to the 4 hydrogen atoms with single bonds.

The formal charge on an atom can be calculated using the following formula: Formal charge = (number of valence electrons) - (number of bonds) - (number of non-bonding electrons) In our case, the formal charge on nitrogen is: Formal charge on N = (5 valence electrons) - (4 bonds) - (0 non-bonding electrons) = +1 The formal charge on each hydrogen atom is: Formal charge on H = (1 valence electron) - (1 bond) - (0 non-bonding electrons) = 0 The overall charge on the \(\mathrm{NH}_{4}+\) ion is: Overall charge = (+1 charge on nitrogen) + (4 hydrogen x 0 charge on each) = +1 This matches the given charge of the ion. Therefore, the Lewis structure for the \(\mathrm{NH}_{4}+\) ion would have a central nitrogen atom connected to four hydrogen atoms, with no remaining lone pairs. The formal charge on the nitrogen atom is +1, and the overall charge of the ion is also +1.