sharing of bonding electrons when covalent chemical bonds are formed electrons (bonding electron...

Sharing of Bonding Electrons

• When covalent chemical bonds are formed electrons (bonding electron pairs) are shared between the bonded atoms. In the simplest cases of homonuclear diatomic molecules the bonding electrons are shared equally. Familiar examples of diatomic molecules are H2, N2, O2, and Cl2 but less familiar examples such as Na2 and K2 are observed in the gas phase.

Sharing of Bonding Electrons

• The bonding electrons in a homonuclear diatomic molecule are equally shared. In the Cl2

molecule the single pair of bonding electrons would be attracted equally to/by each of the two Cl atoms. (One could view the equal sharing of electrons in the Cl2 molecules as a tie in an electrical tug of war – each Cl atom attracting the bonding electron pair to the same extent).

The electrostatic potential maps for sodium chloride, hydrogen chloride and chlorine

FIGURE 10-5Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10 Slide 3 of 48

Sharing of Bonding Electrons – cont’d:

• When heteronuclear diatomics are formed there is no reason to expect the bonding electrons to be equally shared. Why? In our study of the Periodic Table we learned that, for metals, ionization energies are small whereas for nonmetals ionization energies are large. This tells us, again, that metal atoms give up electrons relatively easily whereas nonmetals do not.


• Further, in discussing electron affinities we learned that the addition of an electron to a nonmetal is more energetically favoured than the addition of an electron to a metal. The ionization energy and electron affinity data can be combined to construct a table of dimensionless electronegativity values.


• Electronegativity values (EN’s) can be defined as a linear combination of ionization energies and electron affinities. For element “A”:

• ENA (IA – EAA)• Electronegativity values measure the ability of

an atom in a molecule to attract electrons towards itself. Most electronegative element is F (least electronegative element is Fr). Mention He?

Electronegativity

Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10 Slide 7 of 48

FIGURE 10-6

•Electronegativities of the elements

Electronegativity Values in Practice

• Electronegativity values and electronegativity value differences can be used to predict whether two elements will form:

• (a) Molecules with non-polar covalent bonds. Here the Δ(EN) value is zero. The familiar examples are the homonuclear diatomics.

• (b) Molecules with polar covalent bonds. Here the Δ(EN) value is moderate. Examples are the heteronuclear diatomics CO and ClF.

Electronegativity Values in Practice

• (c) Compounds with ionic bonds. Here the Δ(EN) value is extreme. Examples are NaCl(s), MgO(s) and Mg3N2(s).

• The pure covalent bond (equal electron sharing) and the ionic bond are sometimes regarded as limiting/extreme cases. In practice, chemists often describe bonding in terms of % covalent character and % ionic character – the %’s relate to Δ(EN) values.

Percent ionic character of a chemical bond as a function of electronegativity difference

FIGURE 10-7


Metals/Nonmetal Compounds/Molecules:

• When discussing bonding between metals and nonmetals it’s important to distinguish between a crystalline ionic compound and a gas containing, for example, gaseous AgCl molecules. Metal atoms are easily vaporized using high energy laser pulses

• Mg(s) → Mg(g) Laser Pulse (hν)

Metals/Nonmetal Compounds/Molecules:

• Once in the gas phase the metal atoms can react with nonmetal atoms or molecules to form extremely polar molecules (high % ionic character). The gaseous metal atoms can also be reacted with more complex molecules. Some of the funnier(???) experiments utilize laughing gas (N2O). For example:

• Mg(g) + N2O(g) → MgO(g) + N2(g)

Unstable Molecules

• Aside: Molecules such as MgO(g) are chemically unstable. They can be formed only under unusual conditions – for example, laser ablation and low gas pressures. These interesting species do provide information to test our understanding of chemical bonding. With spectroscopic techniques we can determine the bond distance (internuclear separation) in these molecules and, as well, measure the size of their electric dipole moments.

Further Electron Dot Formulas

• We will use electronegativities in many bonding discussions. We will also use them to predict whether individual bonds in polyatomic molecules are electrically polar. We will also employ electron dot structures (and VSEPR theory) to predict the three dimensional shapes of polyatomic molecules.

10-4 Writing Lewis Structures

• All the valence e- of atoms must appear (ions might have “extra” electrons).

• Usually, the e- are paired.• Usually, each atom requires an octet.

• H only requires 2 e-.• Multiple bonds may be needed.

• Readily formed by C, N, O, S, and P.


Skeletal Structure

• Identify central and terminal atoms.


CH

HH

HCH

H

O

Skeletal Structure• Hydrogen atoms are always terminal atoms (form

one covalent bond).• Central atoms are generally those with the lowest

electronegativity.• Carbon atoms are always central atoms.• Generally structures are compact and symmetrical.


Class Examples – Lewis Structures:

• 1. Draw Lewis structures for CH4, CH2Cl2 and SiH2F2. What do all of these structures have in common? Write formulas for four other molecules with similar Lewis structures. How many chemical bonds are normally formed by C, Si and Ge atoms in a covalently bonded molecule?

• 2. Draw Lewis structures for the OH free radical and the OH- ion.

Problematic Lewis Structures – Formal Charges:

• In some cases more than one plausible Lewis structure can be drawn for a molecule (or polyatomic ion). In such cases, calculations of formal charges for all atoms in a molecule (for all possible Lewis structures) serves to identify the most likely (physically reasonable) Lewis structure.

Formal Charges – Rules:

• Step 1: Determine (for every atom!) how many valence electrons a neutral atom would possess (for example, 6 for oxygen).

• Step 2: Subtract the number of lone pair electrons that a particular proposed Lewis structure assigns to each atom.

• Step 3: Subtract half of the number of bonding electrons surrounding each bonded atom. Why is the factor of ½ needed?

Formal Charge

Copyright © 2011 Pearson Canada Inc. Slide 21 of 48General Chemistry: Chapter 10

FC = #valence e- - #lone pair e- - #bond pair e-2

1

FC(O) = 6 - 4 – (4) = 02

1

FC(N) = 5 - 0 – (8) = +12

1

•• ••

•• ••O=N=O +


Formal Charge of an Alternative Lewis Structure

••

•• ••

•• •• ••O—N—O

FC(O≡) = 6 - 2 – (6) = +12

1

FC(N) = 5 - 0 – (8) = +12

1

FC(O—) = 6 - 6 – (2) = -12

1

••O N O ••

••

••

+ + -


• Sum of FC is the overall charge.• FC should be as small as possible.• Negative FC usually on most electronegative

elements.• FC of same sign on adjacent atoms is unlikely.

General Rules for Formal Charge

+

••O≡N—O ••

••

••

-+

Resonance Structures

• For some molecules (and polyatomic ions) more than one Lewis structure with “identical” patterns of formal charges appears to be plausible for a given molecule (or ion). In such cases we often view the actual molecule to be best represented as a resonance hybrid of the plausible structures. The examples of ozone and sulfur dioxide may be familiar from high school chemistry.

Resonance Structures – Bond Orders:• For O3 and SO2, the two resonance structures

suggest a bond order of 1.5 for each molecule.• The bond order calculation takes the number

of bonds between equivalent atom pairs and divides by the number of bonded atom pairs. In examples such as CO3

2- there are three equivalent atom pairs (and even more for SO4

2- ion).

Resonance Structures – Symmetry:

• Molecules whose structure is best represented as an “average” of several contributing resonance structures often are more symmetrical than an individual Lewis structure might suggest. In ozone, for example, both oxygen-oxygen bond lengths are identical. One “side” of the molecule is a mirror image of the other side.

Resonance


O O O O O O••••

••

••

••••

••••

••••

••

••+ +- -

Electrostatic potential map of ozone

O O O••••

••••

••+ -½ -½

Free Radicals – Unpaired Electrons:

• Many of the most chemically interesting molecules are those with unpaired electrons. Such molecules are necessarily paramagnetic. They are frequently also chemically unstable (reactive and/or short lifetime). The methyl radical, ∙CH3, represented on the next slide is a short lived species. Methyl radicals can (for example) combine to form ethane (stable).

• ∙CH3(g) + ∙CH3(g) → C2H6(g)

Exceptions to the Octet Rule• Odd-Electron Species


••

••

••

•

H—C—H

H

•

•O—H ••

••

N=O

Class Examples

• 1. Draw a Lewis structure for the BF3 molecule. Using all single bonds the molecule appears to be “electron deficient” (less than 8 electrons around B. Can a structure or structures be drawn for which the octet rule is satisfied?

• 2. Draw one or more Lewis structures for the carbonate ion (CO3

2-) for which the octet rule is obeyed. Is the concept of resonance useful?

Incomplete Octets


B

F

FF

••••

••

••

•••• ••••

••

B

F

FF

-

•••• ••

••

••

••

+

••

•• ••

••••

B

F

FF

••

••

•••• ••

•••• -

+

•• ••

“Expanded Octets”

• The next two slides show structures for a molecule containing P and an ion containing S where both P and S have more than the expected octet of “valence shell” electrons.

• What do both of these structures have in common? An extreme cases of hypervalent molecules is IF7 (14 valence shell electrons).

Expanded octets


•• ••

••

P

Cl

ClCl

••••

••

••

••••

•• P

Cl

Cl

••

••Cl

••••

••

••••

•• •• ••••Cl

••••

••Cl•• S

F

F

••

••F

••••

••

••

•••• •• ••

••F

••••••F••

F••

••

••

Expanded Valence Shells


sharing of bonding electrons when covalent chemical bonds are formed electrons (bonding electron...

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