Video by Janet Gray Coonce MS

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How to draw the Lewis dot structure for H20?  We know that it would have to be H-O-H because hydrogen can only form one bond.  Hydrogen is always on the outside of a molecule. Therefore oxygen would have to be in the middle.

To draw the Lewis structure, we first put dots in for the VALENCE electrons.  Valence electrons are the electrons available in the outer shell of an atom.  Hydrogen has only 1 valence electron and gets 1 dot each representing the lone electron associated with the hydrogen protonOxygen has 6 valence electrons.  A full outer shell for oxygen would have 8 valence electrons, 4 groups of 2.  Therefore we draw 2 paired electrons and 2 unpaired electrons.  To form a stable compound, hydrogen would like to have 2 more electrons in it’s outer shell to complete the the octet of 8 electrons.

Both hydrogen and oxygen will be able to have their VALENCE shell of electrons filled if they each SHARE their unpaired electrons with the other.  This way they each will have a pair of electrons in a COVALENT BOND between them.  In the top illustration a line was drawn between the unpaired electron dots.  This was re-drawn for neatness sake in the bottom drawing.  By bonding covalently with another atom, each atom has a full valence shell of electrons and remain electron neutral.

Lewis Structure of ammonia gas, formula NH3

Ammonia solution used as a cleaning solution is ammonia gas NH3 which is dissolved in water.  This can be expressed by the equation NH3 +H20 –> NH4+ +OH.  You are familiar with the pungent odor of ammonia gas which is also used as a respiratory stimulant (smelling salts).

In this case nitrogen is the central atom.  Nitrogen is in group V-A of the periodic table.  It has a total of 5 valence electrons with 3 unpaired electrons which are available for covalent bonding.  Hydrogen is in group IA and has one unpaired electron available for covalent bonding.

Here a line is drawn between the dots representing the electrons involved in each of the covalent bonds.  When these electrons are covalently shared, nitrogen has a total of 8 valence electrons and each hydrogen has 2 valence electrons.  The Lewis dot drawing makes it clear that nitrogen has completed its octet with 3 covalent bonds with hydrogen and 1 lone pair.  Nitrogen now with its full octet of valence electrons is isoelectric with the noble gas neon.  Hydrogen with 2 electrons in its outer shell is isoelectric with the noble gas helium

–Transcription by James C. Gray MD FACOG

by Janet G. Coonce MS

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An ‘s’ orbital is in the shape of a sphere.  A ‘p’ orbital is the shape of a dumbbell.  These shapes define the probability of finding an electron in that space.

So in an s orbital the electron can be found anywhere within a sphere surrounding the nucleus as demonstrated on the left.  In the p orbital the electron can be found anywhere above or below the nucleus within a dumbbell shaped orbital illustrated on the right.

A single bond is a sigma bond.

A double bond is a sigma and a pi bond.

A triple bond is a sigma bond and 2 different pi bonds.

If the orbital overlap occurs between the nuclei as illustrated in the drawings on the left it is referred to as a sigma bond.  All single bonds are sigma bonds.  Pi bonds occur when the overlap occurs above and below the nuclei, but the positively charged nuclei are not directly in line with the overlap area.  This arrangement explains why pi bonds are weaker than sigma bonds.  In the illustration on the right, 2 p orbitals overlap to form a pi bond.  On the left, 3 different sigma bonds are demonstrated.  At the top the sigma bond occurs between 2 s orbitals.  In the middle, it is between an s and a p orbital.  This could also represent an sp hybridized orbital.  In the bottom a single sigma bond is illustrated between 2 p orbitals.  In each of the 3 examples of a sigma bond, the electron overlap occurs between the nuclei of the 2 atoms.

A single bond is a sigma bond.

A double bond is a sigma and a pi bond.

A triple bond is a sigma bond and 2 different pi bonds.

A pi bond is weaker than a sigma bond.

Because a double bond is a sigma and a pi bond, it is stronger than a single bond but not twice as strong.

A triple bond is stronger than a double bond but not 3 times as strong as a single bond.

In this Lewis dot structure of chloromethane (or methyl chloride) each of the 4 single bonds are sigma bonds.

In this Lewis dot structure of ethylene, each carbon atom has 3 sigma bonds.  One sigma bond is between the two carbon atoms and each carbon atom is connected to 2 atoms of hydrogen by 2 additional sigma bonds.  In addition, there is a pi bond between the two carbons ( a double bond consists of a sigma bond and a pi bond).  Each carbon has electrons forming 3 covalent bonds.  Each electron pair within a bond is trying to get as far away from other bonds.  This arrangement is referred to as sp2 hybridized orbitals.  Each bond is equal distant.  Therefore each the 3 bonds is in the same plane (trigonal planar), and each bond angle is 120 (360/3 = 120).

In this 3 dimensional model of ethylene, the yellow play dough represents the dumbbell shaped p orbital which will form the pi bond.  The 3 green orbitals represent the sp2 hybridized orbitals.  The carbon nucleus is at the center of the orbitals.   The hydrogen atoms are represented by the white balls which are bonding with the carbon in the sp2 hybridized orbital.

This illustration shows how the dumbbell shaped p orbitals overlap to form the pi bond and at the same time a sigma bond forms between the green sp2 hybridized orbitals.  All 4 hydrogen atoms are in the same plane.

Triple Bonding in Ethyne (Acetylene)

This is the Lewis dot structure of H2C2 (formula of ethyne better known as acetylene).  Between the carbons is a triple bond consisting of 2 pi bonds and one sigma bond.  There are a total of 6 electrons being shared between carbon atoms across the triple bond.  Between the carbon and hydrogen is an sp hybridized sigma bond which is sharing 2 electrons.  The octet rule is satisfied and each carbon has 8 electrons in the hybridized orbitalsAcetylene has planar geometry because each of the bonds will want to be as far apart as possible.

The above illustration of the 2 sp hybridized orbitals in each carbon in ethyne was taken from the video “Hybridization Geometries & Bond Angles.”  To form the triple bond in ethyne a sigma bond forms between one of the sp hybridized orbitals of each carbon.  Hydrogen forms a sigma bond between the other sp hybridized orbital of each carbon.  The 2 dumbbell shaped p orbitals each form a pi bond.  A triple bond is 2 pi bonds and one sigma bond.

Transcription by James C. Gray MD FACOG

Video by Janet Gray Coonce MS

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Now let’s draw the Lewis dot structure for SO4-2.  That means we will have to add 2 electrons into our Lewis dot structure.  So we have a sulfur atom with 4 oxygen atoms surrounding it.  Oxygen is in period 2 and sulfur is in period 3.  Sulfur is directly below oxygen on the periodic table and they are both in group 6ABoth sulfur and oxygen have 6 valence electrons

We begin drawing the Lewis dot structure by drawing the 6 valence electrons around each sulfur and oxygen atom.

In this step we added an extra electron to each of the 2 oxygen atoms as illustrated.  Now all of the electrons are accounted for and 2 of the oxygen atoms are electro-negative.

In this step we drew a single bond between each atom.  The oxygen atoms on the left and right each have 8 valence electrons and are happy   but rather than forming a double bond, each of them have acquired an extra electron which will satisfy the octet rule but will make the atom electro-negative.  This explains the negative 2 charge of the sulfate ion.  The top and bottom oxygen molecule each have an unpaired electron which is available to join in a bond with sulfur.

Here we draw a bond between the unpaired electrons.  If we count the electrons, each of the oxygen atoms have 8 electrons in their valence shell so the octet rule is satisfied.  Sulfur has 12 electrons.  It can do this.  It is called an expanded octetAtoms belonging to period 3 or more may be able to expand the octet.  Let’s re-draw what we have illustrated here.

T

he two delocalized electrons make multiple resonance structures possible for the sulfate ion (SO4-2). This is the completed Lewis dot structure for one of them.  All 6 of sulfur’s valence electrons are being shared in bonds with oxygen giving a total of 12 valence electrons, an expanded octet.  The sulfur atom is happy with valence shell complete and its formal charge is neutral.

Transcription by James C. Gray MD FACOG

Video by Janet Gray Coonce, MS

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How do we use Lewis dot structures to illustrate bonding in CO2?  We we know that the oxygen atoms will repel each other so carbon will be in the center with an oxygen at each end.  How many valence electrons do each of those atoms have?  Well we know oxygen has six valence electrons because it is in periodic table group 6ACarbon is in group 4A of the periodic table so it has 4 valence electrons.

This Lewis dot structure demonstrates that oxygen has 2 unpaired electrons and has 2 unpaired electrons available for covalent bonds. Carbon has 4 unpaired electrons available for 4 covalent bonds.

The illustration shows how to use a Lewis dot structure to represent the bonding in the CO2 molecule.

Lewis Dot Structure for Methanol

The chemical symbol for methanol is CH30H or CH40.  Hydrogen can only form one bond so it can not be the central atom.  Hydrogen is always on the outside of a molecule with which it is sharing electrons.  We know therefore that Carbon and Oxygen will be in the center with 3 hydrogen atoms surrounding the carbon atom and a single hydrogen bonding the oxygen atom.

Here we demonstrate the valence electrons with dots.  Carbon has 4, hydrogen 1 and oxygen 6.  It is easy to see that when the electrons are shared each atom will have its outer shell completed.  With sharing each hydrogen will have 2 electrons, carbon 8 electrons and oxygen 8 electrons in the outer shell.  The molecule is balanced and happy .

When the dots are connected between the shared electrons, the Lewis dot structure is completed.  This is the Lewis dot structure for methanol

Transcription by James C. Gray MD FACOG

##### Video by Janet Gray Coonce MS

If you plan to view the video on your cell phone, consider your data plan and whether you should wait until you have a WiFi connection to avoid cellular charges.

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Now I’d like to show you how to draw Lewis dot structures, the preferred structure without having to use math.  Sound good?  I like it.  It’s just a visual representation of the valence electrons and how they are going to bond together.  So lets go over guidelines for writing Lewis dot structures.

First we know that HYDROGEN (H) only forms one bond so it is always going to be on the OUTSIDE of a molecule.  It has only 1 valence electron (which is the reason) it is in group 1A of the periodic table

The Lewis dot structure of the hydrogen atom is shown here (H-).  It is always going to want to form one bond so that this hydrogen atom will feel like it has 2 electrons (a pair) by sharing an electron with another atom, say a hydrogen atom H–H.  When two atoms of hydrogen share an electron in a bond to form a molecule of hydrogen gas, each hydrogen atom is happy .  By sharing its valence electron covalently with each other, each atom in the H2molecule has its duet (a pair) of electrons in its valence shell.  This diatomic molecule of hydrogen gas (H2) is sharing a pair of electrons in a covalent bond and is represented by the Lewis structure (H–H).  The H2 molecule is isoelectronic to helium which is a noble gas.

CARBON has a valence of 4.  It always wants to form 4 bonds and we can remember this by looking at the Lewis dot structure for carbon.  Carbon has 4 valence electrons.

Lewis dot structures are drawn in a fashion to illustrate the octet rule:  atoms of low (<20) atomic number tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas.

If carbon picked up 4 more electrons by sharing in a bond then the outer shell would be filled, the octet rule would be satisfied and carbon would be happy .  So carbon wants to form 4 bonds.  It can be in the form of 4 single bonds; or a double bond and 2 single bonds; or a single bond and a triple bond.  The Triple bond is where a total of 6 electrons are shared between carbon and another atom.  When 2 more electrons are shared with another atom in a single bond, carbon’s outer shell is full (8 e) and carbon is happy .  When the octet rule is satisfied the carbon atom is isoelectronic with the noble gas neon.  The compounds you will find in nature that contain carbon are called organic compounds.

Most often carbon forms either 4 single bonds; or a double bond and 2 single bonds; or a single bond and a triple bond as illustrated in these Lewis dot structures.  In each situation illustrated, there are a total of 4 bonds around the carbon atom.  There are a total of 4 pairs of electrons (8 total, an octet) in its valence shell.  There are exceptions to this rule and they are called carbocations and carbanions but lets not talk about those now.  That’s for organic chemistry later.

The next element you need to know is OXYGEN.  What is its Lewis dot structure?  We know it has 6 valence electrons because it is in the periodic table Group 6A

The Lewis dot structure for oxygen has 6 electrons drawn around the oxygen symbol, 2 LONE PAIRS and 2 electrons available for bonding as illustrated.  Oxygen wants to pick up 2 electrons so when each lone electron is covalently bonded to another atom, the octet rule will be satisfied.  The 2 bonds may either be 2 single bonds or a double bond.  Either way when each of the lone electrons are shared, there will be 8 electrons in its outer shell.  The outer shell is full, isoelectronic with the noble gas neon, and this makes it happy .  Oxygen in our atmosphere exists as the diatomic molecule O2.

In summary: Hydrogen has a valence of 1 and it will always want to form a bond outside the atom.  Hydrogen can only form 1 bond.

Carbon has 4 valence electrons in its outer shell and will want to form 4 bonds.

Oxygen has 6 valence electrons in its outer shell and will want to form 2 bonds and there will be 2 lone pairs.

Now lets look at the element NITROGEN

Nitrogen has the atomic symbol N.  It is in periodic table Group 5A so there are 5 valence electrons in its outer shell.  The Lewis dot structure has 5 electrons drawn around it.  It has one lone pair and 3 valence electrons available for bonding.  If it can pick up 3 electrons, it will be able to complete the octet rule which will make it stable.

Nitrogen bonds are represented by 3 single bonds and a lone pair;  or a single bond, a double bond and a lone pair;  or a lone pair and a triple bond.  In each case, the octet rule is satisfied and nitrogen is isoelectric with the noble gas neon.  Like hydrogen and oxygen, elemental nitrogen gas exists as a diatomic molecule.

Now lets look at the HALOGENS.  Halogens belong to Group 7A of the periodic table.  The halogens include fluorine (F), chlorine (C), bromine (B), and iodine (I).  They all have 7 valence electrons.  Therefore In a Lewis dot structure, any of the halogens can be represented by the letter X surrounded by 7 electrons represented by 3 lone pairs and a single electron available for bonding.  The halogen atom wants to pick up an electron.  It can do that through an ionic bond where the electron is completely transferred which will give the anion a negative charge, or it can share with another atom in a covalent bond.  When it shares a covalent bond it “feels” like it has the 8 valence electrons needed to complete the outer shell.

The halogen is represented by 3 lone pairs and one electron available for bonding.

Transcribed by James C. Gray, MD FACOG