15 Lewis Structures

(1)

Lewis Dot Structures

(2)

Molecular Structure & Bonding

A molecular structure, unlike a simple

molecular formula, indicates the exact 3-D nature of the molecule. It indicates which

(3)

Molecular Formula vs. Molecular

Structure

Molecular formula – H₂O

Molecular structure: _{.. ..}

O

(4)

Molecular Structure

Two issues:

(5)

What is stuck to what?

The first thing you need to do in drawing a molecular structure is to figure out which

atom sticks to which other atoms to generate a skeletal model of the molecule.

(6)

Lewis Dot Structures

The first step towards establishing the full 3-D geometry of a molecule is determining what is stuck to what and how each atom is

connected.

(7)

Two Rules

1. Total # of valence electrons – the total

number of valence electrons must be accounted for, no extras, none missing.

2. Octet Rule – every atom should have an

(8)

What’s a “valence electron”?

It’s an electron in the outermost shell of an atom. When two atoms bump, it’s the valence shells that hit first.

(9)

The Bohr Model

Nucleus p

p n p _n

n n

(10)

e-Electronic Structure of Atoms

Since the electrons are so important,

understanding the electronic structure of

atoms is critical to understanding why atoms react with each other.

(11)

What’s wrong with the Bohr

Model?

Nucleus p

p n p _n

n n

(12)

It doesn’t



The electron “orbits” are stable and electrons can move between them by absorbing light (higher energy orbitals) or emitting light

(13)

Quantum electronic structure

The solution to the electron paradox is that the world of the atom is not “classical” but

“quantum mechanical”.

In a quantum world, only certain discrete

(14)

Can’t do it! Has to jump!

Nucleus p

p n p _n

n n

(15)

e-Can’t do it! Has to jump!

Nucleus p

p n p _n

n n

(16)

e-Electron Orbitals

 Electron orbitals are diffuse. The electron is not a

hard little pellet, but a “probability cloud”.

 Electron orbitals are 95% probability intervals.

 Allowed electron orbitals are determined by 4

(17)

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Electron Orbitals

Every electron is represented by 4 quantum numbers. These electron Quantum numbers are:

n = principal quantum number (kind of like the Bohr orbit)

l = angular momentum quantum number – gives the shape

m_l = magnetic quantum number – determines the number of orbitals of a given shape (2l+1)

(20)

Allowed Quantum numbers

n = 1, 2, 3, 4, 5….

l = 0, 1, 2, 3, 4…(n-1)

There are as many different types of orbitals as “n”.

m_l = -l, -l+1…-1, 0, 1…l-1, l

There are 2l+1 orbitals of a given type (l)

m_s = -1/2, 1/2

(21)

Possible Quantum numbers

(22)

Possible Quantum numbers

n = 1 l=0 m_l = 0 m_s=-1/2

m_s=+1/2 n = 2 l=0 m_l = 0 m_s=-1/2

m_s=+1/2 l=1 m_l = -1 m_s=-1/2

m_s=+1/2 m_l = 0 m_s=-1/2

m_s=+1/2 m_l = 1 m_s=-1/2

(23)

What do these numbers mean?

n is like the Bohr orbit number. It gives the “shell” the electron is in.

l is the orbital number, it specifies the type of orbital within the same shell.

m_l gives the orientation of the orbital – these are different flavors of the same orbital

(24)

Shorthand Notation

Orbitals are specified by letters:

l=0 is an s orbital

l=1 is a p orbital

l=2 is a d orbital

l=3 is an f orbital

(25)

Shorthand notation

n=1, l=0 is called a 1s orbital n=2, l=0 is called a 2s

n=2, l=1 is called a 2p n=3, l=2 is called a 3d

The number of electrons in each orbital are indicated as a superscript.

1s2 means 2 electrons are in the 1s orbital

(26)

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Rules Governing Electrons

1. Pauli Exclusion Principle - No two electrons

in an atom can have the same 4 quantum numbers

2. Lowest energy orbitals fill first

3. Hund’s rule – Electrons pair up as a last

resort

4. An orbital being full or half-full is good!

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Energy of the Orbitals

1s

2s 2p

3s 3p 3d 4s 4p 4d 4f

5s 5p 5d 5f 5g

(29)

Pauli Exclusion Principle

This determines the number of orbitals in a shell and the total number of electrons that fit in each orbit.

There is only 1 orbital (s) in the 1st shell of only 1 type which can hold, at most, 2 electrons.

There are 2 different orbitals (s, p) in the 2nd_{shell. There is 1}

type of s (always) and 3 types of p (always). Each type can hold 2 electrons. So, at most, the 2nd shell can hold 8

electrons.

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Electron Configurations

If you need to figure out the electron

configuration, you just count the electrons

and start filling from lowest energy to highest.

For example, consider C

(31)

Carbon

C – 6 electrons

1s is the lowest energy orbital, it takes 2 2s is the next lowest, it also takes 2

2p comes next, it can take up to 6, so it gets the last 2 electrons

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Mg = 12 electrons

1s gets 2 2s gets 2 2p gets 6 3s gets 2

(35)

Clicker question

What is the ground state electron configuration of N?

A. 1s22s5

B. 1s22s22p3

C. 1s22s22p5

(36)

Clicker question

What is the ground state electron configuration of As?

A. 1s22s22p63s23p83d104p3

B. 1s22s22p63s23p64s24p3

C. 1s22s22p63s23p64s23d104p3

(37)

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Clicker question

What is the ground state electron configuration of Cr?

A. 1s22s22p63s23p83d4

B. 1s22s22p63s23p64s23d4

C. 1s22s22p63s23p64s13d5

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Core vs. Valence Electrons

Core electrons – completed shells

Valence electrons – “outer” or incomplete shells

(41)

What’s the valence configuration

of…

Ca?

[Ar]4s2

Mo?

[Kr]5s14d5

Ga?

(42)

(43)

What about Fe

3+

?

Fe (atomic number 26) [Ar]4s23d6

Take away 3 electrons… [Ar]4s23d3

OR

(44)

Ions are different…

Electrons go into the shells in the order we indicated, but to form an ion by removing

electrons, they come out in a different order. N=4 electrons come out before N=3. N=5

electrons come out before N=4.

(45)

What about Fe

3+

?

Fe (atomic number 26) [Ar]4s23d6

Take away 3 electrons…the 4s electrons come out before 3d, so…

[Ar]4s23d3 - NOT

(46)

Molecular Structure & Bonding

A molecular structure, unlike a simple

molecular formula, indicates the exact 3-D nature of the molecule. It indicates which

(47)

Molecular Formula vs. Molecular

Structure

Molecular formula – H₂O

Molecular structure: _{.. ..}

O

(48)

Molecular Structure

Two issues:

(49)

What is stuck to what?

The first thing you need to do in drawing a molecular structure is to figure out which

atom sticks to which other atoms to generate a skeletal model of the molecule.

(50)

Lewis Dot Structures

The first step towards establishing the full 3-D geometry of a molecule is determining what is stuck to what and how each atom is

connected.

(51)

Two Rules

1. Total # of valence electrons – the total

number of valence electrons must be accounted for, no extras, none missing.

2. Octet Rule – every atom should have an

(52)

Determining the number of

valence electrons:

Full d-orbitals do not count as valence

electrons. They belong to the inner shell.

For example:

As is [Ar]4s23d104p3

(53)

How many valence electrons

does Ge have?

(54)

(55)

Take a look at Ge electron

structure

[Ar]4s23d104p2

(56)

How many valence electrons

does Ti have?

(57)

How many valence electrons

does Te have?

(58)

Total Number of Valence Electrons

The total number of available valence electrons is just the sum of the number of valence electrons that

each atom possesses (ignoring d-orbital electrons)

So, for H₂O, the total number of valence electrons = 2 x 1 (each H is 1s1) + 6 (O is 2s22p4) = 8

(59)

Central Atom

In a molecule, there are only 2 types of atoms:

1. “central” – bonded to more than one other atom. 2. “terminal” – bonded to only one other atom.

(60)

How many central atoms in

ethanol?

(61)

Bonds

Bonds are pairs of shared electrons.

Each bond has 2 electrons in it.

You can have multiple bonds between the same 2 atoms. For example:

C-O C=O C O

(62)

Lewis Dot Structure

Each electron is represented by a dot in the structure

.

:Cl: ¨

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Drawing Lewis Dot Structures

1. Determine the total number of valence

electrons.

2. Determine which atom is the “central” atom. 3. Stick everything to the central atom using a

(64)

Dot structure for H

₂

O

1. Total number of valence electrons: 6 + (2 x 1) =8

2. Central Atom – typically, the central atom will be

leftmost and/or bottommost in the periodic table. It is the atom that wants more than one thing stuck to it. H is NEVER the central atom.

(65)

Dot structure for H

₂

O

(66)

Drawing Lewis Dot Structures

1. Determine the total number of valence electrons. 2. Determine which atom is the “central” atom.

3. Stick everything to the central atom using a single

bond.

4. Fill the octet of every atom by adding dots.

5. Verify the total number of valence electrons in the

(67)

Dot structure for H

₂

O

_..

H – O – H ¨

That is a total of 8 valence electrons used:

(68)

1. Determine the total number of valence electrons.

2. Determine which atom is the “central” atom.

3. Stick everything to the central atom using a single bond.

5. Verify the total number of valence electrons in the structure.

6. Add or subtract electrons to the structure by making/breaking

bonds to get the correct # of valence electrons.

(69)

Formal Charge of an atom

“Formal charge” isn’t a real charge. It’s a pseudo-charge on a single atom.

Formal charge = number of valence electrons – number of bonds – number of non-bonding electrons.

Formal charge (FC) is ideally 0, acceptably +/-1, on occasion +/- 2. The more 0s in a structure, the better.

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Dot structure for H

₂

O

_..

H – O – H ¨

FC (H) = 1-1-0 = 0

FC (O) = 6 – 2 – 4 = 0

(71)

Another example

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(74)

CO

₂

CO₂

Total number of valence electrons = 4 from carbon + 2x6 from oxygen = 16

Central Atom?

(75)

CO

₂

CO₂

16 total valence electrons

O – C – O

Fill out the octets

.. .. ..

(76)

Drawing Lewis Dot Structures

(77)

CO

₂

16 total valence electrons

.. .. .. :O – C - O: ¨ ¨ ¨

Structure has 20 electrons in it. Too many!

I need to lose 4 electrons. What’s the best way to do that?

(78)

CO

₂

:O = C = O: ¨ ¨

Structure has 16 electrons in it. Just right!

Notice, this works because there are 2 ways to count the electrons:

1. When I count the total # of electrons, I count each

electron once.

2. When I count the electrons for each atom, I count

(79)

CO

₂

:O = C = O: ¨ ¨

Is this the only structure I could have drawn?

I only needed two new bonds, I didn’t specify where they needed to go!

.. :O C - O: ¨ ..

:O - C O: ¨

(80)

Choosing between different

structures?

The first test is formal charge: :O = C = O:

¨ ¨

FC (O) = 6 – 2 – 4 = 0 FC (C) = 4 – 4 – 0 = 0 ..

:O C - O: ¨

FC (left O) = 6 – 3 – 2 = 1 FC (C) = 4 – 4 – 0 = 0

FC (right O) = 6 – 1 – 6 = -1

(81)

Are these even different?

.. :O C - O: ¨ ..

:O - C O: ¨

(82)

Are they different?

.. :O₁ C – O₂ : ¨ ..

:O₁ - C O₂ : ¨

If I label them, I can see a difference. (Isotopic labeling).

(83)

Resonance

.. :O₁ C – O₂ : ¨ ..

:O₁ - C O₂ : ¨

Structures that are identical, but differ only in the arrangement of bonds are called resonance structures.

(84)

Resonance

When you have resonance, the real structure is not any one of the individual structures but

the combination of all of them.

You can always recognize resonance – there are double or triple bonds involved.

(85)

Resonance

Resonance is indicated by drawing all resonance structures, separated by “ ”

.. ..

:O C - O: :O - C O: :O = C = O: ¨ ¨ ¨ ¨

But this is not necessary in this case, as the last

(86)

Nitrite ion

Draw the Lewis Dot structure for NO₂

-How many valence electrons?

N has 5, O has 6, but there’s one extra (it’s an ion!)

(87)

Nitrite LDS

What’s the central atom?

Nitrogen O – N – O .. .. .. :O – N - O: ¨ ¨ ¨

(88)

Nitrite LDS

.. .. .. :O – N - O: ¨ ¨ ¨

How do you fix the problem? Make a bond

.. .. .. :O = N - O: ¨

(89)

Nitrite LDS

.. .. .. .. .. .. :O = N - O: :O - N = O: ¨ ¨

What’s the real structure look like?

It’s an average of those 2. Kind of 1-1/2 bonds

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Double bond between C and O

or C and Cl?

A.. C and O B. C and Cl

C. Doesn’t matter

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2-Exceptions to the Octet Rule

There are exceptions to the octet rule:

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Incomplete Octets

The most common elements that show incomplete octets are B, Be besides H.

So, for example, BCl₃ has the Lewis structure:

.. .. : Cl – B – Cl: ¨ | ¨ : Cl : ¨

Total valence electrons is correct at 24. FC (B) = 3 - 3 – 0 = 0

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Expanded Octets

The most common atoms to show expanded octets are P and S. It is also possible for some transition metals.

An example of an expanded octet would be PCl₅:

.. ..

:Cl: :Cl: Total valence e- = 40 .. ..

:Cl – P - Cl : FC(P) = 5 – 5 – 0 =0 ¨ | ¨