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Lewis Dot Structures: Resonance 7:01 Covalent Bonds: Predicting Bond Polarity and Ionic Character 4:56 VSEPR Theory & Molecule Shapes 9:09 Dipoles & Dipole Moments.

• Lewis Dot Structure Generator Nh3
• Lewis Dot Structure Generator N2
• Lewis Dot Structure Generator Ch4
• Lewis Dot Structure Generator So2
• Lewis Dot Structure Generator Free

Learning Objectives

• Draw Lewis structures for covalent compounds.

The following procedure can be used to construct Lewis electron structures for more complex molecules and ions.

How-to: Constructing Lewis electron structures

1. Determine the total number of valence electrons in the molecule or ion.

• Lewis structure of sulfate ion is drawn in this tutorial step by step. Total valence electrons concept is used to draw the lewis structure of SO 4 2-. In lewis structure of sulfate ion, there should be charges on several atoms due to -2 charge.
• Lewis Dot Structures for Covalent Compounds Part 1 CLEAR & SIMPLE. Bonding Models and Lewis Structures Lewis Structure Generator.

• Add together the valence electrons from each atom. (Recall that the number of valence electrons is indicated by the position of the element in the periodic table.)
• If the species is a polyatomic ion, remember to add or subtract the number of electrons necessary to give the total charge on the ion.

For CO₃²⁻, for example, we add two electrons to the total because of the −2 charge.

2. Arrange the atoms to show specific connections.

• When there is a central atom, it is usually the least electronegative element in the compound. Chemists usually list this central atom first in the chemical formula (as in CCl₄ and CO₃²⁻, which both have C as the central atom), which is another clue to the compound’s structure.
• Hydrogen and the halogens are almost always connected to only one other atom, so they are usuallyterminal rather than central.

3. Place a bonding pair of electrons between each pair of adjacent atoms to give a single bond.

• In H₂O, for example, there is a bonding pair of electrons between oxygen and each hydrogen.

4. Beginning with the terminal atoms, add enough electrons to each atom to give each atom an octet (two for hydrogen).

• These electrons will usually be lone pairs.

5. If any electrons are left over, place them on the central atom.

• We will explain later that some atoms are able to accommodate more than eight electrons.

6. If the central atom has fewer electrons than an octet, use lone pairs from terminal atoms to form multiple (double or triple) bonds to the central atom to achieve an octet.

• This will not change the number of electrons on the terminal atoms.

7. Final check

• Always make sure all valence electrons are accounted for and that each atom has an octet of electrons, except for hydrogen (with two electrons).
• The central atom is usually the least electronegative element in the molecule or ion; hydrogen and the halogens are usually terminal.

Now let’s apply this procedure to some particular compounds, beginning with one we have already discussed.

Example (PageIndex{1}): Water

Write the Lewis Structure for H₂O.

Solution

Example (PageIndex{2})

Write the Lewis structure for the (CH_2O) molecule

Solution

Exercise (PageIndex{1})

Write Lewis electron structures for CO₂ and SCl₂, a vile-smelling, unstable red liquid that is used in the manufacture of rubber.

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The United States Supreme Court has the unenviable task of deciding what the law is. This responsibility can be a major challenge when there is no clear principle involved or where there is a new situation not encountered before. Chemistry faces the same challenge in extending basic concepts to fit a new situation. Drawing of Lewis structures for polyatomic ions uses the same approach, but tweaks the process a little to fit a somewhat different set of circumstances.

Writing Lewis Structures for Polyatomic Ions

Recall that a polyatomic ion is a group of atoms that are covalently bonded together and which carry an overall electrical charge. The ammonium ion, (ce{NH_4^+}), is formed when a hydrogen ion (left( ce{H^+} right)) attaches to the lone pair of an ammonia (left( ce{NH_3} right)) molecule in a coordinate covalent bond.

When drawing the Lewis structure of a polyatomic ion, the charge of the ion is reflected in the number of total valence electrons in the structure. In the case of the ammonium ion:

(1 : ce{N}) atom (= 5) valence electrons

(4 : ce{H}) atoms (= 4 times 1 = 4) valence electrons

subtract 1 electron for the (1+) charge of the ion

total of 8 valence electrons in the ion

It is customary to put the Lewis structure of a polyatomic ion into a large set of brackets, with the charge of the ion as a superscript outside of the brackets.

Exercise (PageIndex{2})

Draw the Lewis electron dot structure for the sulfate ion.

Answer

Exceptions to the Octet Rule

As important and useful as the octet rule is in chemical bonding, there are some well-known violations. This does not mean that the octet rule is useless—quite the contrary. As with many rules, there are exceptions, or violations.

There are three violations to the octet rule. Odd-electron molecules represent the first violation to the octet rule. Although they are few, some stable compounds have an odd number of electrons in their valence shells. With an odd number of electrons, at least one atom in the molecule will have to violate the octet rule. Examples of stable odd-electron molecules are NO, NO₂, and ClO₂. The Lewis electron dot diagram for NO is as follows:

Although the O atom has an octet of electrons, the N atom has only seven electrons in its valence shell. Although NO is a stable compound, it is very chemically reactive, as are most other odd-electron compounds.

Electron-deficient molecules represent the second violation to the octet rule. These stable compounds have less than eight electrons around an atom in the molecule. The most common examples are the covalent compounds of beryllium and boron. For example, beryllium can form two covalent bonds, resulting in only four electrons in its valence shell:

Boron commonly makes only three covalent bonds, resulting in only six valence electrons around the B atom. A well-known example is BF₃:

The third violation to the octet rule is found in those compounds with more than eight electrons assigned to their valence shell. These are called expanded valence shell molecules. Such compounds are formed only by central atoms in the third row of the periodic table or beyond that have empty d orbitals in their valence shells that can participate in covalent bonding. One such compound is PF₅. The only reasonable Lewis electron dot diagram for this compound has the P atom making five covalent bonds:

Formally, the P atom has 10 electrons in its valence shell.

Example (PageIndex{3}): Octet Violations

Identify each violation to the octet rule by drawing a Lewis electron dot diagram.

1. ClO
2. SF₆

Solution

a. With one Cl atom and one O atom, this molecule has 6 + 7 = 13 valence electrons, so it is an odd-electron molecule. A Lewis electron dot diagram for this molecule is as follows:

b. In SF₆, the central S atom makes six covalent bonds to the six surrounding F atoms, so it is an expanded valence shell molecule. Its Lewis electron dot diagram is as follows:

Exercise (PageIndex{3}): Xenon Difluoride

Identify the violation to the octet rule in XeF₂ by drawing a Lewis electron dot diagram.

The Xe atom has an expanded valence shell with more than eight electrons around it.

Summary

Lewis dot symbols provide a simple rationalization of why elements form compounds with the observed stoichiometries. A plot of the overall energy of a covalent bond as a function of internuclear distance is identical to a plot of an ionic pair because both result from attractive and repulsive forces between charged entities. In Lewis electron structures, we encounter bonding pairs, which are shared by two atoms, and lone pairs, which are not shared between atoms. Lewis structures for polyatomic ions follow the same rules as those for other covalent compounds. There are three violations to the octet rule: odd-electron molecules, electron-deficient molecules, and expanded valence shell molecules.

Contributions & Attributions

This page was constructed from content via the following contributor(s) and edited (topically or extensively) by the LibreTexts development team to meet platform style, presentation, and quality:

• Modified by Joshua Halpern (Howard University)

• CK-12 Foundation by Sharon Bewick, Richard Parsons, Therese Forsythe, Shonna Robinson, and Jean Dupon.

• Marisa Alviar-Agnew (Sacramento City College)

• Henry Agnew (UC Davis)

Lewis structure of sulfate ion is drawn in this tutorial step by step. Total valence electrons concept is used to draw the lewis structure of SO₄^2-. In lewis structure of sulfate ion, there should be charges on several atoms due to -2 charge.

Sulfate ion | sulfate ion | SO₄^2-

Sulfate ion is one of the oxyanion of sulfur. Sulfur is at +6 oxidation state in SO₄^2-. Also, sulfate ion has a -2 charge.

Lewis structure of SO₄^2-

There are two S=O bonds and two S-O bonds in sulfate ion lewis structure. Sulfur atom is the center atom and four oxygen atoms are located around sulfur atom. There are no lone pairs in the last shell of sulfur atom.

Steps of drawing lewis structure of SO₄^2-

Following steps are required to draw the SO₄^2- lewis structure and they are explained in detail in this tutorial.

1. Find total number of electrons of the valance shells of sulfur and oxygen atoms
2. Total electrons pairs
3. Center atom selection
4. Put lone pairs on atoms
5. Check the stability and minimize charges on atoms by converting lone pairs to bonds.

Drawing correct lewis structure is important to draw resonance structures correctly

Total number of electrons of the valance shells of SO₄^2-

Both Sulfur and oxygen atoms are located at VIA group in the periodic table. So, oxygen and sulfur atoms have six electrons in their valence shell.

• Total valence electrons given by sulfur atom = 6

There are four oxygen atoms in SO₄^2- ion, Therefore

• Total valence electrons given by oxygen atoms = 6 *4 = 24

There are -2 charge on SO₄^2- ion. Therefore there are two more electrons which comes from outside to contribute to the total valence electrons.

• Total valence electrons = 6 + 24 + 2 = 32

Total valence electrons pairs

Total valance electrons pairs = σ bonds + π bonds + lone pairs at valence shells

Total electron pairs are determined by dividing the number total valence electrons by two. For, SO₄^2- ion, Total pairs of electrons are 16.

Center atom of SO₃^2- ion

To be the center atom, ability of having greater valance is important. Therefore sulfur has the more chance to be the center atom (See the figure) because sulfur can show valance of 6. Maximum valnce of oxygen is two. So, now we can build a sketch of SO₄^2- ion.

Lone pairs on atoms

• There are already four S-O bonds in the above sketch. Therefore only twelve (16-4 = 12) valence electrons pairs are remaining.
• First, mark those twelve valence electrons pairs as lone pairs on outside atoms (on oxygen atoms). One oxygen atom will take three lone pairs following the octal rule (oxygen atom cannot keep more than eight electrons in its valence shell).
• For four oxygen atoms, twelve electrons pairs are spent. Now all electron pairs are spent. There is no electron pairs to mark on sulfur atom.

Charges on atoms

After, marking electron pairs on atoms, we should mark charges of each atom. Each oxygen atom will get a -1 charge and sulfur atom get a +2 charge. The overall charge of ion is ( -1*4 + (+2) ) = -2.

Check the stability and minimize charges on atoms by converting lone pairs to bonds

Lewis Dot Structure Generator Nh3

When charges exist everywhere (on atoms) in the ion or molecule, that structure is not stable. We should try to reduce charges on atoms as much as possible. Now, we are going to learn how these facts will affect on sulfate ion.

• Oxygen atoms should hold negative charges because electronegativity of oxygen is higher than sulfur. Otherwise, we can say, ability of holding negative charges is great in oxygen atoms than sulfur atoms.
• The drawn structure is not a stable one because all atoms have charges.
• Now, we should try to minimize charges by converting lone pair or pairs to bonds. So convert one lone pair of one oxygen atom to make a new S-O bond.
• Now there is a double bond between sulfur atom and one oxygen atom. Now, there are three S-O single bonds between sulfur atom and other three oxygen atoms.

You see charges of atoms are reduced. Now, there is no charge in one oxygen atom and charge of sulfur atom is reduced from +2 to +1. So we have an stable ion than out previous one.

Can I reduce charges of atoms furthermore?

You should know, sulfur can keep more than eight electrons in its last shell. Therefore we can convert one more lone pair of another oxygen atom to a bond.

In new structure, charges of atoms are reduced than previous structure. Now there are no any charge on sulfur atom and two oxygen atom. Also, only two oxygen atoms have -1 negative charges. Now you understand this structure of SO₄^2- is more stable than previous structures. So, this structure has more chance to be the lewis structure of SO₄^2- ion.

Lewis Dot Structure Generator N2

Lewis structure of SO₄^2- (sulfate) ion

Lewis Dot Structure Generator Ch4

Questions

the total number of electrons in the valence shells of all atoms in the ion SO₄^2- is?

There are 32 electrons in valence shells of all atoms in the ion SO₄^2- ion.

Lewis Dot Structure Generator So2

Lewis Dot Structure Generator Free

Is number of lone pairs similar in S₂O₃3^2- and SO₄3^2- lewis structures?

Yes. Compare both lewis structures. Otherwise, we can think an oxygen atom of sulfate ion is replaced by a sulfur atom. Because both sulfur and oxygen belongs to group 6, their valence electrons in last shell is similar.