1. Covalent Bonding The Nature of Covalent Bonding Bonding Theories
Polar Bonds and
Molecules

2. Ch. 16.1 The Nature of Covalent Bonding
Covalent bonds occur between atoms that are “sharing” electrons
Form covalent compounds
There is a kind of “tug of war” for the electrons
There can be single, double and triple covalent bonds
Single bond – a bond in which 2 atoms share a pair of electrons
Double bond – bond that involves 2 shared pairs of e-
Triple bond – bond that involves 3 shared pairs of e-

3. Ch. 16.1 The Nature of Covalent Bonding
When writing formulas for covalent bonds:
Pairs of electrons are represented by a dash
Called structural formulas
Chemical formulas that show the arrangement of atoms in molecules and polyatomic ions
Each dash between atoms in structural formulas indicates a pair of electrons
Dashes are never used to show ionic bonds
Chemical formulas
Chemical formulas of ionic cmpds describe formula units
Chemical formulas of covalent cmpds describe molecules

4. Ch. 16.1 The Nature of Covalent Bonding
Combinations of atoms of non-metallic atoms are likely to form covalent bonds
Groups 4A, 5A, 6A, and 7A
Summarized by G. Lewis in the octet rule
sharing of e- occurs if atoms achieve noble gas config.
H2 is an exception to this rule

5. Ch. 16.1 The Nature of Covalent Bonding
Group trends
Halogens form single covalent bonds in their diatomic molecules (ex: F – F )
Chalcogens form double covalent bonds in their diatomic molecules (ex: O = O)
Phicogens form triple covalent bonds in their diatomic molecules ( N = N )
The Carbon group tends to form 4 bonds with other atoms

6. Ch. 16.1 The Nature of Covalent Bonding
Covalent bonding can be explained using electron configurations and orbital boxes
Show H2O, NH3 and CH4 on board
Double and triple covalent bonds
Oxygen forms a double bond in a diatomic molecule
It is an exception to the octet rule, 2 unpaired e- (p. 442)
Nitrogen forms a triple bond in a diatomic molecule
Satisfies the octet rule, all e- are paired
Multiple covalent bonds can form between unlike atoms (ex: CO2, CH3OH)

7. Ch. 16.1 The Nature of Covalent Bonding
Coordinate covalent bonds
Carbon monoxide (CO) is a covalent bond that cannot be explained by “normal” covalent bonding (p. 444)
When C and O form a traditional double bond, the O atom satisfies the octet rule, but the C atom does not
For both atoms to achieve a stable octet, oxygen donates one of its unshared pairs of e- for bonding
Called a coordinate covalent bond
The structural formula shows an arrow pointing from the atom donating the e- pair to the atom receiving the e- pair
Once formed, it is like any other covalent bond

8. Ch. 16.1 The Nature of Covalent Bonding
Coordinate covalent bonds
Most polyatomic cations and anions contain covalent and coordinate covalent bonds
The polyatomic ammonium ion (NH4+) has a coordinate covalent bond
Polyatomic ions contain both covalent and ionic bonds (see pg. 446 H3O+)

9. Ch. 16.1 The Nature of Covalent Bonding
Bond dissociation energies
The release of heat in a bond suggests the product is more stable that the reactants
The total energy required to break the bond between covalently bonded atoms is called the bond dissociation energy
The higher the dissociation energy, the stronger the energy

10. Ch. 16.1 The Nature of Covalent Bonding
Resonance
Resonance structures are those that occur when it is possible to write 2 or more valid electron dot formulas that have the number of electron pairs for a molecule or ion
Scientists once believed that the e- pairs flipped back and forth (resonated) between structures
The actual bonding is a hybrid of the extremes represented by the resonance structures
Double-headed arrows are used to connect resonance structures (see pg O3)

11. Ch. 16.1 The Nature of Covalent Bonding
Exceptions to the octet rule
For some molecules or ions, it is impossible to write structures that satisfy the octet rule
This occurs when the total number of valence e- is an odd number
NO2 is an example (see pg. 450) – an unpaired e- exists in each of the resonance structures

12. Ch. 16.1 The Nature of Covalent Bonding
Magnetism and bonding
Electrons are spinning charged particles with a magnetic field
Paired e-, spinning in opposite directions, cancel out the magnetic field (diamagnetic substances have all electrons paired)
These are weakly repelled by an external magnetic field
Unpaired e- result in an external magnetic field (paramagnetic substances contain 1 or more unpaired e-)
These are strongly attracted to an external magnetic filed

13. Ch. 16.1 The Nature of Covalent Bonding
Ferromagnetism is the strong attraction of Fe, Co, and Ni to magnetic fields
These ions all have unpaired electrons
Electrons line up in magnetic domains, remain in that state permanently
Covalent bond lengths
Single bonds are longest, double bonds are intermediate, and triple bonds are the shortest
Exceptions to the octet rule
There are many molecules that do not follow the octet rule, yet exist as stable bonds
A few atoms (esp. P and S) sometimes expand the octet to include 10 or 12 electrons (see pg. 451)

14. Ch. 16.2 Bonding Theories Molecular orbitals
When 2 atoms combine, their atomic orbitals overlap to produce molecular orbitals (orbitals that apply to the entire molecule)
The molecular orbital model of bonding requires that the number of molecular orbitals equal the number of overlapping atomic orbitals
When 2 atomic orbitals overlap, 2 molecular orbitals are created
One is called a bonding orbital, the other is called an anti- bonding orbital
The anti-bonding orbital has a higher energy that the atomic orbitals from which it formed

15. Ch. 16.2 Bonding Theories Molecular orbitals
When H2 forms, the 1s atomic orbitals overlap
2 electrons are available for bonding (see pg. 452)
The energy of the e- in the bonding molecular orbital is lower than the e- in the atomic orbitals of the separate H atoms
Electrons seek the lowest energy level, so they fill the bonding molecular orbital
This makes a stable covalent bond between the H atoms
The anti-bonding orbital is empty
Sigma and pi bonds are caused by the overlapping of “s” and “p” orbitals

16. Ch. 16.2 Bonding Theories VSEPR Theory
The valence-shell electron-pair repulsion theory
States that because e- pairs repel, molecular shape adjusts so that the valence-electron pairs are as far apart as possible
Unshared pairs are important when you are trying to predict the shapes of molecules
Unshared pairs are held more closely to the atom
Unshared pairs strongly repel bonding pairs of e-, pushing them closer together than might be expected

17. Ch. 16.2 Bonding Theories Molecular geometries Tetrahedral Pyramidal
all 4 bond angles are 109.5o
CH4 is an example
Pyramidal
bond angles between shared pairs is 107o
NH3 is an example
Bent
bond angles are 105o
H2O is an example

18. Ch. 16.2 Bonding Theories Molecular geometries Trigonal planar Linear
Bond angles are 120o
An example is BF3
Linear
Bond angles are 180o
CO2 is an example
Trigonal bipyramidal
Bond angles of 120o and 90o
An example is PF5

19. Ch. 16.2 Bonding Theories Hybrid orbitals
In orbital hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals
One 2s and three 2p orbitals of a carbon atom overlap to form an sp3 hybrid orbital
These are at the tetrahedral angle of 109.5o
Four sp3 orbitals of carbon overlap with the 1s orbitals of the four hydrogen atoms
This allows for a great deal of overlap, which results in the formation of 4 C-H sigma bonds
These are unusually strong covalent bonds

20. Ch. 16.3 Polar Bonds and Molecules
Bond polarity
Covalent bonds involve the sharing of electrons
However, they can differ in how the bonds are shared
Depends on the kind and number of atoms joined together
When electrons are shared equally, a nonpolar covalent bond is formed
When the atoms share the electron unequally, a polar covalent bond is formed
The more electronegative element will have the stronger electron attraction and will acquire a slightly negative charge
The less electronegative element will acquire a slightly positive charge

21. Ch. 16.3 Polar Bonds and Molecules
Bond polarity
The lower case Greek letter delta is used to show partial positive and negative charges (p.461)
Table 16.4 (p. 462) shows how the size of the electronegativity difference between bonded atoms can predict the type of bond that will form
These are the most probable bonds that will form
The boundary between covalent and ionic is not always distinct
Polar molecules have one end that is slightly positive and one end that is slightly negative
A molecule with 2 different poles is called a dipole

22. Ch. 16.3 Polar Bonds and Molecules
Attractions between molecules
Molecules are attracted to each other by a variety of forces
These are intermolecular forces
They are weaker than either ionic or covalent bonds
They are largely responsible for determining whether a substance is a solid, liquid or gas
There are several kinds of intermolecular forces
Van der Waals forces
Dispersion forces
Dipole interactions
Hydrogen bonds

23. Ch. 16.3 Polar Bonds and Molecules
Intermolecular forces
Van der Waals forces
The weakest attractions are known collectively as van der Waals forced
Named after Johannes van der Waals
Consist of two possible types
Dispersion forces
The weakest of all intermolecular forces
Caused by the motion of electrons
Strength increases as the number of electrons increases
Dipole interactions
Occurs when polar molecules are attracted to each other
Slightly negative regions are attracted to slightly positive ones

24. Ch. 16.3 Polar Bonds and Molecules
Intermolecular forces
Hydrogen bonds
Attractive forces in which a hydrogen covalently bonded to a to a very electronegative atom is also weakly bonded to to an unshared pair of another electronegative atom
Always involves hydrogen
The strongest of all the intermolecular forces
Extremely important in determining the properties of water and biological molecules such as proteins

25. Ch. 16.3 Polar Bonds and Molecules
Intermolecular attractions and molecular properties
The physical properties of a compound depend on the type of bonding it has
whether the bonding is ionic or covalent
Table 16.5 (p. 465)
Lists the characteristics of ionic and covalent compounds
Melting and boiling points of most covalent compounds are low compared to ionic compounds
A few molecular solids are very stable substances - network solids in which all atoms are covalently bonded to one another
An example is diamond
These compounds have extremely high melting and boiling points, or may simply vaporize above a certain temperature