From 2fa8b010b0164f8f3b3bb87bf53743469f7a2c48 Mon Sep 17 00:00:00 2001 From: eggy Date: Tue, 13 Oct 2020 19:45:42 -0400 Subject: [PATCH] chem: useful tip for lewis structure confusion --- docs/sch3uz.md | 2 +- 1 file changed, 1 insertion(+), 1 deletion(-) diff --git a/docs/sch3uz.md b/docs/sch3uz.md index 5d32e6c..d340d95 100644 --- a/docs/sch3uz.md +++ b/docs/sch3uz.md @@ -313,7 +313,7 @@ Sometimes, one atom in a covalent bond may contribute both electrons in a shared ### Formal charge -There may be several correct ways to draw covalent bonds in Lewis structures and dot diagrams. Solving for the **formal charge** of each atom involved in a covalent bond can help identify the **best** structure to construct. The formal charge of an atom in a covalent bond represents the charge that that atom has. The sum of all formal charges in a covalently bonded compound is equal to the charge of the overall compound. +There may be several correct ways to draw covalent bonds in Lewis structures and dot diagrams. Solving for the **formal charge** of each atom involved in a covalent bond can help identify the **best** structure to construct. The formal charge of an atom in a covalent bond represents the charge that that atom has. The sum of all formal charges in a covalently bonded compound is equal to the charge of the overall compound. **The element with the lowest electronegativity is almost always in the centre.** The formal charge of an atom can be calculated using the following equation: $$\text{Formal charge} = \text{# of valence electrons of element} - \text{# of unpaired electrons} - \text{# of covalent bonds}$$