1. Chapter 7 Covalent Bonding and Electron-Dot Structures

2. The Covalent Bond
Bonds: a force that holds groups of two or more atoms together and makes them function as a unit
Required 2 e- to make a bond

3. Every bond in every molecule has its own specific bond length, the distance between nuclei to two bonded atoms. However, the magnitudes of the various attractive and repulsive forces between nuclei and electrons in a covalent bond depend on how close the atoms are.
If the hydrogens atoms are too far apart, the attractive forces are small and no bond exists. IF they are too close together, the repulsive interaction between the nuclei becomes so strong that it pushes the atom part.
Therefore, bond length is measured at the optimum distance between nuclei where net attractive forces are maximized and H—H molecule is most stable.

4. Lines between the atoms indicate shared electrons in covalent molecules
Every covalent bond has its own characteristic length that leads to maximum stability and that is roughly predictable from the knowledge of atomic radii.
- The bond length of diatomic halogen increases down the group as the radius of element increases.
Chapter 7, Unnumbered Figure 2, Page 225

5. Strengths of Covalent Bonds
The amount of energy that must be supplied to break a chemical bond in an isolated molecule in the gaseous state is equivalent to the amount of energy released when the bond forms. It is called bond dissociation energy (D).
Chapter 7, Unnumbered Figure 1, Page 225

6. Periodic properties and bond length can be used to explain why some bonds are stronger than other.
As we proceed down through the period of halogen, the value of D becomes smaller, meaning the bond is weaker. As the atomic radius of the halogen increases, the shared electrons are farther away and more shielded from the positive charge nucleus, leading to a longer and weaker bond.
There are exceptions, shorter bonds are typically stronger. For example, F—F bond is predicted to be stronger than Cl—Cl bond because fluorine atoms are smaller, but in fact the bond dissociation energy for F—F is 159 kJ/mol compared to 243 kJ/mol for Cl—Cl. (Ch.8 discussion)

7. The correlation between bond length and strength can be explain thru bond order, referring to the number of electron pairs shared between atoms.
Multiple bonds are both shorter and stronger than their corresponding single-bond because there are more shared electrons holding atoms together.
Bond order =
Bond order =

8. Polar Covalent Bonds: Electronegativity
Chapter 7: Covalent Bonds and Molecular Structure
2/17/2019
Polar Covalent Bonds: Electronegativity
Depending on the relative electronegativities of the two atoms sharing electrons, there may be partial transfer of electron density from one atom to the other. When the electronegativities are not equal, electrons are not shared equally and partial ionic charges develop.
The extend of electron transfer in a compound is most easily visualized with what called electrostatic potential maps, which use color to portray the calculated electron distribution in an isolated, gas-phase molecule.
Cl2
NaCl
HCl
Copyright © 2008 Pearson Prentice Hall, Inc.

9. Polarity
Polar covalent bonds – the bonding electrons are attracted somewhat more strongly by one atom in a bond
Electrons are not completely transferred
More electronegative atom: δ- . (δ represents the partial negative charge formed)
Less electronegative atom: δ+

10. Polar Covalent Bonds: Electronegativity
American chemist, Linus Pauling developed the electronegativity scale call Electronegativity. It is a measure of the ability of an atom in a molecule to draw bonding electrons density to itself.
Atoms with greater electronegativities will attract more of the shared electron density to themselves, causing a “polarity” to the bond.
Fluorine, the most highly electronegative element. Metallic elements on the left o the periodic table attract electrons only weakly and are the least electronegative elements. Halogens and other reactive nonmetals in the upper right of the table attract electrons strongly and are most electronegative.

11. Relationship Between Electronegativity and Bond Type
Polarity of a bond is another important factor influencing bond dissociation energy.
The attraction between the partial charges increases the energy required to break the bond.
Increasing bond polarity leads to increased bond strength.
Predicting bond polarity
Atoms with similar electronegativity (Δ EN <0.4) –form nonpolar bond
Atoms whose electronegativity differ by more than two (Δ EN > 2) – form ionic bonds
Atoms whose electronegativity differ by less than two (Δ EN < 2) – form polar covalent bonds

12. Polar bond vs. Ionic bond
Chapter 7, Unnumbered Figure 1, Page 229

13. Example
For each of the following pairs of bonds, choose the bond that will be more polar
a. H-P, H-C b. N-O, S-O
An electrostatic potential map of water is shown. Which atom, H or O, is positive (electron-poor) and which is negative (electron-rich?

14. Lewis Structures or Lewis Formula
represents how an atom’s valence electrons are distributed in a molecule
Show the bonding involves (the maximum bonds can be made)
Try to achieve the noble gas electron configuration
The common pattern

15. Examples
Show the short-hand electron configuration then draw the Dot Lewis structure for the following atoms: Na C Mg Cl

16. Rules for Wring Dot Lewis structure
Step 1: Calculate the total number of valence electrons of all atoms in the molecule
Step 2: Create a skeletal structure using the following rules:
Hydrogen atoms (if present) are always on the “outside” of the structure. They form only one bond
The central atom is usually least electronegative. It is also often unique (i.e,. the only one atom of the element in the molecule). Remember, there might be no “central” atom.
Connect bonded atoms by line (2-electron, covalent bonds
Step 3: Place lone pairs around outer atoms (except hydrogen) so that each atom has an octet
Step 4: Calculate the number of electrons you haven’t used. Subtract the number of electrons used so far, including electrons in lone pair and bonding pairs, from the total in Step 1. Assign any remaining electrons to the central atom as lone pair
Step 5: If the central atom is B (boron) or Be (beryllium), skip this step
If the central atom has an octet after step 4, skip this step
If the central atom has only 6 electrons, move a lone pair from an outer atom to form a double bond between outer atom and the central atom
If the central atom has only 4 electrons, do Step 5a to two different outer atoms (i.e, form two double bonds) or twice to one outer atom (i.e., form one triple bond)

18. Rules for drawing Dot Lewis structure
Duet Rule: sharing of 2 electrons
E.g H2
H : H
Octet Rule: sharing of 8 electrons
Carbon, oxygen, nitrogen and fluorine always obey this rule in a stable molecule
E.g F2, O2
Bonding pair: two of which are shared with other atoms
Lone pair or nonbonding pair: those that are not used for bonding

19. Electron-Dot Structures
Chapter 7: Covalent Bonds and Molecular Structure
2/17/2019
Electron-Dot Structures
Think of this section as an introduction. It is much easier to write electron-dot structures using the rules listed in the next section.
Copyright © 2008 Pearson Prentice Hall, Inc.

20. Drawing Lewis formula with one central atom
Draw the Lewis structure for the following
HBr NH3
H2O CCl4

21. Lewis formula for ions
When calculating the sum of valence electrons for all atoms, add one additional for each negative charge in anion and subtract one electron for each positive charge in cation.
NH CO3-2

22. Lewis formula for multiple bonds
If the central atom has only 6 electrons, move a lone pair from an outer atom to form a double bond between outer atom and the central atom
If the central atom has only 4 electrons, do Step 5a to two different outer atoms (i.e, form two double bonds) or twice to one outer atom (i.e., form one triple bond)
CH2O HCN

23. Lewis Formula with Expanded Octet
Elements with empty d orbitals can have more than 8 electrons
SF POCl3

24. Electron-dot structures of compounds containing only hydrogen and second-row elements
Molecular that contain carbon atoms bonded together in a chain are called organic compounds.
CH2F2 CH3CH3  

25. Rules for drawing resonance
when there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures
the actual molecule is a combination of the resonance forms – a resonance hybrid
it does not resonate between the two forms, though we often draw it that way
There are a number of common misconceptions that students have about resonance structures:
Resonance structures must have the same connectivity
Resonance structures must have the same number of electrons
Only electron positions can change (BUT electrons do not jump back and forth), they are delocalized.
Never use an equilibrium arrow

26. Resonance
Lewis structures are supposed to represent the arrangement of valence electrons in a molecule, including predictions of bond order and bond length. However, in some cases, one single Lewis structure doesn't do an adequate job of representing (modeling) the actual arrangement of valence electrons in a molecule
For example, for the case of ozone, O3. Experiments have proven that ozone is a symmetric molecule with two identical oxygen-oxygen bond distances
However, a simple Lewis structure for ozone would predict that it should have one long (_______________) bond and one short (________________) bond. The lack of agreement between experiment and theory indicates that there is a deficiency in the model. In this case, the problem with the model is that it assumes that a pair of electrons can only be shared between two nuclei—which is not always true.

27. Resonance Calculate bond order in ozone Bond order =
Ozone doesn’t have one O=O double bond and one O—O single bond as the individual structure imply; rather ozone has two equivalence O—O that we can think of having a bond order of 1.5, midway between pure single bonds and pure double bonds (as shown in the calculation above). Both bonds have an identical length of 128pm.
The actual molecule is intermediate between two (or more) resonance contributors is called a resonance hybrid.
The concept of resonance is an adaptation of the Lewis formula that helps account for complexity of actual molecules (more can be learned in Organic Chemistry).

28. Example
The nitrate ion, NO3-, has three equivalence oxygen atoms, and its electronic structure is a resonance hybrid of three electron-dot structure. Draw them and predict the bond order

29. Example
Called “laughing gas”, nitrous oxide (N2O) is sometimes used by dentists as an anesthetic. Given the connections N-N-O, draw two electrons-dot resonance structures for N2O.