Lewis Dot Structure Covalent Bonds Calculator

1. Lewis Dot Structure Covalent Bonds
2. Lewis Dot Structure For Covalent Compounds
3. Lewis Dot Structure Covalent Bonds Calculator Formula

Could someone explain the lewis structure diagram of covalent compound Al2Cl6?

Get the free 'Lewis structure' widget for your website, blog, Wordpress, Blogger, or iGoogle. Find more Chemistry widgets in Wolfram Alpha. Using Lewis Dot Symbols to Describe Covalent Bonding The valence electron configurations of the constituent atoms of a covalent compound are important factors in determining its structure, stoichiometry, and properties. For example, chlorine, with seven valence electrons, is one electron short of an octet.

Information: Steps for Drawing Lewis Structures for Covalent Compounds Study the two examples in the table of how to write structures for CO3 2-and NH3. Make sure you understand each of the five steps. CO3 2-NH 3 Step #1: Add up the number of valence electrons that should be included in the Lewis Structure. 4 + 3(6) + 2 = 24 (carbon has four. Electrons in the Lewis Dot Structure? ZDraw Lewis Structures for O 2 and N 2. 14 Covalent Bond zThere is an optimum distance between atoms in a covalent bond. This is the bond length: calculated by adding the radii of two atoms. Strengths of Covalent Bonds zCloser nuclei result in a stronger bond. Shorter bond = stronger bond. 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.

Also, why is Al2Cl6 (aluminium chloride) covalent?

Here's what we have to know for school but I don't know how it works. What do the arrows mean?

1 Answer

Lewis Dot Structure Covalent Bonds

Explanation:

Chlorine has 17 electrons, but 10 of those are in the orbitals of the lower energy levels. (#1S, 2S, 2P# orbitals).

Lewis Dot Structure For Covalent Compounds

These are completely enveloped by the larger #3S#-orbital (think of a golfball inside a tennisball) so takes no part in the formation of bonds.

The other seven are distributed in the #3S# and the three #3P#-orbitals, but upon forming (Covalent) bonds these orbitals hybridise into #SP#-orbitals
for more info about Hybridisation, here's a good link:

In our case the #3S#- and three #3P#-orbitals will hybridise into #SP^3#-orbitals.:

(Picture courtesy of https://en.wikipedia.org/wiki/Orbital_hybridisation)

Each orbital can contain, and indeed strives to contain, 2 electrons (#e^-#).

Chlorine thus has 7 electrons in the 4 #SP^3#-orbitals: 3 orbitals are filled with 2 #e^-# each, the fourth has only one. If you look the the picture below (that I copied from Above) you will see 6 electrons paired in 3 'filled orbitals'

The fourth one, the one that contains the single, unpaired #e^-#, joins in the fourth #SP^3#-orbital with one from the Al-atom (the one on the right). So in this bond between the Al-atom and the Cl-atom, each donates a single electron. That's why the bond is represented by a straight line.

Aluminium has only 3 electrons: 2 in the #3S# and one in one of the #3P#-orbitals. However, upon hybridisation these 3 electrons are spread over 4 #SP^3#-orbitals. Like in the one mentioned above, the other two cooperate as well in the formation of covalent bonds with 2 Chlorine atoms These are circled in Green:

Chlorine has a rather high Electronegativity, which means that it pulls rather hard at electrons from other atoms (from each other, and from other elements.

It is a tug-of-war that the Aluminium atom is threatening to lose, leaving it rather #delta^+# (positive).

At the same time, the Chlorine atoms are satisfied in their hunger for electrons, in fact they don't want any more because they have a full set of 8!
Because of this 'diminished appetite' on the side of the Chlorine atoms, and the increased hunger on the part of the Aluminium atom, The bond is formed by BOTH of the Chlorine electrons.
This is the explanation for the arrow circled by Aqua:

In #Al_2Cl_6#, this happens twice but I'm sure you can spot the other one by now?

Hope this helps...

PS: By the way, it is covalent because the bonds are created by sharing of electrons between the two atoms.

Related questions

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.

Lewis Dot Structure Covalent Bonds Calculator Formula

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

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