Lewis Structures (The Localized Electron Model) G. N. Lewis 1875 - 1946

Using electron-dot symbols, G. N. Lewis developed the Localized Electron Model of chemical bonding (1916) in which valence electrons exist as lone pairs or as individual electrons seeking to form a pairing in order to achieve an octet. Later, Linus Pauling would expand the Localized Electron Model to include Resonance and Orbital Hybridization, Collectively known as Valence Bond Theory (1930). In 1957 VSEPR Theory was added to predict molecular geometry, also describing any resulting molecular polarity in molecules.

According to Lewis Theory, there are two types of valence electrons: •Non-bonding (or unshared) pairs

Localized Electron Model In Lewis’s Localized Electron Model, molecules are described as being composed of atoms that are bound together by sharing pairs of electrons. He was able to show that the arrangement of atoms in Linus and Ava Helen Pauling in Munich, molecules could be predicted based on with Walter Heitler the arrangements of valence electrons (left) and Fritz London (right). 1927 of all atoms involved in the molecule. Walter Heitler and Fritz London (1927) were the first to solidify Lewis’s idea by linking atomic orbital overlap to Schrödinger’s wave equation (1925) to show how two hydrogen atom wave functions join together to form a covalent bond.

We have seen how we can build models of molecules by combining atoms according to electron dot structures. ..

:Br:

•Bonding single (or unpaired) electrons

+ 3 •Boron has three unpaired electrons therefore it can form three covalent bonds •Bromine has three unshared pairs and one unpaired electron, therefore it can only form one covalent bond. •What about nitrogen?

=

.. : N Br .. : :Br .. :

Today, we are going to learn a process by which we will be able to draw a model of any molecule.

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Lewis Structures

Writing Lewis Structures

PCl3 5 + 3(7) = 26

Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

Writing Lewis Structures

Writing Lewis Structures • Things to consider when building primary skeleton:

2. Build a reasonable skeletal structure for the molecule using only single bonds. Keep track of the electrons: 26 − 6 = 20


The central atom should be the least electronegative element that isn’t hydrogen.

Writing Lewis Structures

Keep track of the electrons: 26 − 6 = 20 − 18 = 2

3. Subtract the total number of electrons used in the primary bonds from the available valence electrons. 4. Fill the octets of the outer atoms by adding unshared pairs

1. Find the sum of valence electrons of all atoms in the molecule from the group number or electron dot structure.

Keep track of the electrons: 26 − 6 = 20


Oxygen never bonds to itself, except in O2 and O3
Carbon atoms are usually bonded to each other
In molecules containing both H and O, hydrogen is usually bonded to oxygen
Carbons should always be saturated with hydrogens when possible

Writing Lewis Structures 5. Fill the octet of the central atom. 6. Check to see that all atoms have and octet and that the correct number of Keep track of the electrons: valence electrons were used 26 − 6 = 20 − 18 = 2 − 2 = 0

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Writing Lewis Structures

Writing Lewis Structures

7. If you run out of electrons before the central atom has an octet… Example: Try building a Lewis structure for HCN

1. Let’s try drawing the Lewis Structures for the following molecules: A. Carbon tetrachloride B. Ammonia C. Oxygen D. Carbon dioxide E. Dihydrogen carbon monoxide F.

Ethanal (C2H4O)

2. Let’s draw the Lewis structure for dihydrogen sulfate and for the sulfate anion formed when dihydrogen sulfate is placed in water.

5. …form multiple bonds until it does.

•Polyatomic ions are formed from a class of molecules called Acids, or in some rare cases, from Bases. •Polyatomic ions are formed as acids or bases loose or gain hydrogen atoms. For example: Hydrogen nitrate

NO3nitrate ion

Hydrogen nitrate looses a hydrogen proton when placed in water, resulting in the formation of the nitrate ion (notice the 1- charge)

•Lewis structures for polyatomic ions must account for the loss or gain of valence electrons
Cations – decrease valence electrons by amount of charge
Anions – increase valence electrons by amount of charge •Lewis structures for polyatomic ions are written in brackets [ ] with the charge denoted as a superscript.

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3.Try drawing the Lewis structure for hydrogen nitrate and the nitrate ion.

4. Draw the Lewis Structure for ozone, O3.

You may notice more than one Lewis structure can be drawn for these species.

Notice that two L.S. can be drawn correctly for ozone, O3 RESONANCE Resonance theory, developed by Lewis (1928), is a key component of valence bond theory and arises when no single conventional model using only even number of electrons shared exclusively by two atoms can actually represent the observed molecule. Resonance involves modeling the structure of a molecule as an intermediate, or average, between several simpler but incorrect structures.

Resonance • But this is at odds with the true, observed structure of ozone, in which…
…both O—O bonds are the same length.

Resonance • One Lewis structure cannot accurately depict a molecule such as ozone. • We use multiple structures, resonance structures, to describe the molecule. • Resonance is denoted by a double headed arrow separating the different Lewis Structures:

Resonance • In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. • They are not localized, but rather are delocalized.

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Resonance

Resonance • Observe HCO2- :

Just as green is a synthesis of blue and yellow…

…ozone is a synthesis of these two resonance structures.

• In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. • They are not localized, but rather are delocalized.

5. Draw all three resonance structures for the nitrate ion.

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