Covalent Lewis Dot Structure Calculator

For example, consider the Lewis dot structure for carbon dioxide. This is a linear molecule, containing two polar carbon-oxygen double bonds. However, since the polar bonds are pointing exactly 180° away from each other, the bond polarities cancel out, and the molecule is nonpolar. Drawing the Lewis Structure for Cl 2 CO. Viewing Notes: The Lewis structure for Cl 2 CO requires you to place Carbon in the center of the structure since it is the most electronegative. You'll need a double bond between the Carbon and Oxygen atoms to acheive full outer shells for the atoms while still only using 24 valence electrons. Get the free 'Lewis structure' widget for your website, blog, Wordpress, Blogger, or iGoogle. Find more Chemistry widgets in Wolfram Alpha. In this interactive and animated object, students distribute the valence electrons in simple covalent molecules with one central atom. Six rules are followed to show the bonding and nonbonding electrons in Lewis dot structures. The process is well illustrated with eight worked examples and two interactive practice problems. Every chemistry student has to learn how to draw Lewis Dot Structures. The key is to understand the steps and practice. Lewis Structures are important to learn because they help us predict: the shape of a molecule. How the molecule might react with other molecules. The physical properties of the molecule (like boiling point, surface tension, etc.).

Covalent Lewis Dot Structures

A bond is the sharing of 2 electrons.

Covalent bonds share electrons in order to form a stable octet around each atom in the molecules. Hydrogen is the exception it only requires 2 electrons (a duet) to be stable.

How do we draw a covalent Lewis Dot Structure?

Level 1 (basic)

1. Add up all the valance electrons of the atoms involved. ex CF4

So C has 4 and F has 7 (x4 we have 4Fs) = 32 valence electrons

2. You need to pick the central atom. This is usually easy, this atom will be surrounded by the others. Never H.

So C will be surrounded by F's.

3. Now we create our skeleton structure by placing bonds in. A bond is a dash that represents 2 electrons.

We have now placed 8 electrons as 4 bonds. We have 32-8= 24 more to place.

4. Starting with the outer atoms add the remaining electrons in pairs until all the electrons have run out.

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All 32 electrons are now in place, count the dots around each F. 6 dots and a bond (2 electrons) is 8. We have our octet.

The carbon has 4 bonds (2electrons) for its 8.

DONE

Level 2 (Double and Triple bonds)

Same rules apply until #4

1. Add up all the valance electrons of the atoms involved. ex CO2

So C has 4 and O has 6 (x2 ) = 16 valence electrons

2. You need to pick the central atom. This is usually easy, this atom will be surrounded by the others. Never H.

So C will be surrounded by O's.

3. Now we create our skeleton structure by placing bonds in. A bond is a dash that represents 2 electrons.

We have now placed 4 electrons as 2 bonds. We have 16-4=12 more to place.

4. Starting with the outer atoms add the remaining electrons in pairs until all the electrons have run out.

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All 16 electrons are now in place, count the dots around each O. 6 dots and a bond (2 electrons) is 8. We have our octet.

The carbon has 2 bonds (2electrons) for its 4....?

We need 8, so move a pair of electrons from the O to between the C and O. It will share 2 pairs of electrons instead of 1. It now has a double bond instead of a single bond.

carbon has 6 electrons, so move 2 from the other oxygen

now they all have an octet, it cleans up like this

Make it symmetrical.

Level 3-Lewis Dots of Polyatomic Ions

Same rules apply, at the end they get brackets and a charge

AP Chemistry and or College Level Rules

1. Determine whether the compound is covalent or ionic. If covalent, treat the entire molecule. If ionic, treat each ion separately. Compounds of low electronegativity metals with high electronegativity nonmetals (DEN > 1.7) are ionic as are compounds of metals with polyatomic anions. For a monoatomic ion, the electronic configuration of the ion represents the correct Lewis structure. For compounds containing complex ions, you must learn to recognize the formulas of cations and anions.

2. Determine the total number of valence electrons available to the molecule or ion by:

(a) summing the valence electrons of all the atoms in the unit and
(b) adding one electron for each net negative charge or subtracting one electron for each net positive charge. Then divide the total number of available electrons by 2 to obtain the number of electron pairs (E.P.) available.

3. Organize the atoms so there is a central atom (usually the least electronegative) surrounded by ligand (outer) atoms. Hydrogen is never the central atom.

4. Determine a provisional electron distribution by arranging the electron pairs (E.P.) in the following manner until all available pairs have been distributed:

a) One pair between the central atom and each ligand atom.
b) Three more pairs on each outer atom (except hydrogen, which has no additional pairs), yielding 4 E.P. (i.e., an octet) around each ligand atom when the bonding pair is included in the count.
c) Remaining electron pairs (if any) on the central atom.

5. Calculate the formal charge (F) on the central atom.

a) Count the electrons shared as bonds. Total = b
b) Count the electrons owned as lone pairs. Total = n
c) F = V - (n + b/2), where V = number of valence electrons for the atom.

6. If the central atom formal charge is zero or is equal to the charge on the species, the provisional electron distribution from (4) is correct. Calculate the formal charge of the ligand atoms to complete the Lewis structure.

7. If the structure is not correct, calculate the formal charge on each of the ligand atoms. Then to obtain the correct structure, form a multiple bond by sharing an electron pair from the ligand atom that has the most negative formal charge.

a) For a central atom from the second (n = 2) row of the periodic table continue this process sequentially until the central atom has 4 E.P. (an octet).
b) For all other elements, continue this process sequentially until the formal charge on the central atom is reduced to zero or two double bonds are formed.

8. Recalculate the formal charge of each atom to complete the Lewis structure.

on to Formal Charge

Chemical Demonstration Videos

Covalent Lewis Dot Structure Calculator

In 1916, ten years before the Schrodinger wave equation, G. N. Lewis suggested that a chemical bond involved sharing of electrons. He described what he called the cubical atom, because a cube has 8 corners, to represent the outer valence shell electrons which can be shared to create a bond. This was his octet rule.

  1. Count the number of valence e- each atom brings into the molecule.For ions, the charge must be taken into account.

    How many valence electrons in BeCl2?

    How many valence electrons in NO2- and NO2+?

  2. Put electron pairs about each atom such that there are 8 electrons around each atom (octet rule), with the exception of H, which is only surrounded by 2 electrons. Sometimes it's necessary to form double and triple bonds. Only C, N, O, P and S (rarely Cl) will form multiple bonds.

    Draw the Lewis dot structure for CF4.

    The number of valence electrons is 4 + 4 ( 7 ) = 32 electrons.

    So, we obtain:

    Draw the Lewis dot structure for CO.

    The number of valence electrons is 4 + 6 = 10 electrons or 5 pairs. Since both C and O allow multiple bonds we can still follow the octet and write:

  3. If there is not enough electrons to follow the octet rule, then the least electronegative atom is left short of electrons.

    Draw the Lewis dot structure for BeF2.

    In BeF2 number of valence e- = 2+ 2(7) = 16 e- or 8 pairs. Since neither Be or F form multiple bonds readily and Be is least electronegative we obtain:

  4. If there are too many electrons to follow the octet rule, then the extra electrons are placed on the central atom.

    Draw the Lewis dot structure for SF4.

    In SF4 the number of valence electrons is 6 + 4 ( 7 ) = 34 electrons or 17 pairs. Placing the extra electrons on S we obtain:

How can the octet rule be violated in this last example? The octet rule arises because the s and p orbitals can take on up to 8 electrons. However, once we reach the third row of elements in the periodic table we also have d-orbitals, and these orbitals help take the extra electrons. Note that you still need to know how the atoms are connected in a polyatomic molecule before using the Lewis-Dot structure rules.

Covalent Lewis Dot Structure

Covalent lewis dot structure

Lewis Dot Structure Practice Worksheet

Homework from Chemisty, The Central Science, 10th Ed.

Lewis Dot Diagram Calc

8.45, 8.47, 8.49, 8.51, 8.53, 8.55, 8.57, 8.59, 8.61, 8.63