Chapter 8 Covalent Bonding

Bond Type Using Electronegativities Nonpolar (0 – 0.3) Covalent Polar (0.3 – 1.7) Ionic Ionic (1.7 – 3.3) Covalent Bonds atoms are joined by sharing of electrons. Nonpolar-covalent: an equal sharing of electrons/an even distribution of electron clouds. Polar-covalent: an unequal sharing of electrons/an uneven distribution of electron clouds.

“molecule”: a neutral group of atoms held together by covalent bonds. Ionic Bonds: a bond that results from a mutual electrical attraction between a cation and an anion. Na + Cl “molecule”: a neutral group of atoms held together by covalent bonds. “diatomic molecule”: a molecule containing only two atoms. ex: O2, H2, N2

Electron Dot Diagrams shows how an atom’s valence electrons become involved in bonding. Hydrogen (1 valence electron) Beryllium (2 valence electron) Nitrogen (5 valence electron) Carbon (4 valence electron) Oxygen (6 valence electron) Neon (8 valence electron)

Bonding site Unshared Pairs/ Lone Pairs Unshared pairs (lone pairs): a pair of electrons that are not involved in bonding (exclusive only to that atom). Lewis Structures Shows how atoms are bonded in a molecule. Draw the dot diagrams for each element first, then “connect the dots”.

Ex: H2O Ex: BF3 Ex: C2H4

Atoms of metals lose valence electrons when forming an ionic compound. Atoms of nonmetals gain electrons. Enough atoms of each element must be used in the formula so that electrons lost equal electrons gained.

Solve Revision of ionic bond formation 2 a. Start with the atoms. • • • K and O • • • •

Solve Apply the concepts to this problem. 2 a. In order to have a completely filled valence shell, the oxygen atom must gain two electrons. These electrons come from two potassium atoms, each of which loses one electron. K • • • O + K+ 2–

Bonding and Interactions The electrostatic forces between the oppositely charged ions hold the cations and anions together in an ionic compound. Ionic compounds generally have high melting points and can conduct an electric current in solution and in the molten state.

O O B Cl Molecular Shapes VSEPR Theory (Valence Shell Electron Pair Repulsion) pairs of electrons are arranged as far apart from each other as possible. Linear: 180o between bonds Ex: O2 Linear O O 2. Trigonal Planar: “flat triangle” – one central atom surrounded by three atoms with NO LONE PAIRS. Ex: BCl3 B Cl Trigonal Planar

C H N H 3. Tetrahedral: four atoms bonded to a central atom. 4. Trigonal pyramidal: three atoms bonded to a central atom with one lone pair of electrons about the central atom. N H

5. Bent: a central atom bonded to two other atoms AND one or two LONE PAIRS. Hybridization see overheads

Molecular Polarity Polar molecules occur because of their uneven distribution of charge. The negative end will be at the more electronegative atom. If bond polarities extend equally and symmetrically in different directions, then the molecule as a whole will be nonpolar polarity of a molecule can be determined by the arrangement of the atoms in a molecule, better known as……………. MOLECULAR SHAPE!

Nonpolar Linear Polar Polar Bent Nonpolar Polar Polar Pyramidal Atoms are Identical Not identical Nonpolar Polar Bent Polar Trigonal Planar Atoms are Identical Not identical Nonpolar Polar Pyramidal Polar Tetrahedral Atoms are Identical Not identical Nonpolar Polar

Hybridization: The blending of Orbitals + = Poodle + Cocker Spaniel = Cockapoo + = s orbital + p orbital = sp orbital

What Proof Exists for Hybridization? We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds. Lets look at a molecule of methane, CH4. Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it.

Carbon ground state configuration What is the expected orbital notation of carbon in its ground state? Can you see a problem with this? (Hint: How many unpaired electrons does this carbon atom have available for bonding?)

Carbon’s Bonding Problem You should conclude that carbon only has TWO electrons available for bonding. That is not not enough! How does carbon overcome this problem so that it may form four bonds?

Carbon’s Empty Orbital The first thought that chemists had was that carbon promotes one of its 2s electrons… …to the empty 2p orbital.

However, they quickly recognized a problem with such an arrangement… Three of the carbon-hydrogen bonds would involve an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.

This would mean that three of the bonds in a methane molecule would be identical, because they would involve electron pairs of equal energy. But what about the fourth bond…?

The fourth bond is between a 2s electron from the carbon and the lone 1s hydrogen electron. Such a bond would have slightly less energy than the other bonds in a methane molecule.

This bond would be slightly different in character than the other three bonds in methane. This difference would be measurable to a chemist by determining the bond length and bond energy. But is this what they observe?

The simple answer is, “No”. Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane. Chemists have proposed an explanation – they call it Hybridization. Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy.

In the case of methane, they call the hybridization sp3, meaning that an s orbital is combined with three p orbitals to create four equal hybrid orbitals. These new orbitals have slightly MORE energy than the 2s orbital… … and slightly LESS energy than the 2p orbitals.

sp3 Hybrid Orbitals Here is another way to look at the sp3 hybridization and energy profile…

sp Hybrid Orbitals While sp3 is the hybridization observed in methane, there are other types of hybridization that atoms undergo. These include sp hybridization, in which one s orbital combines with a single p orbital. This produces two hybrid orbitals, while leaving two normal p orbitals

sp2 Hybrid Orbitals Another hybrid is the sp2, which combines two orbitals from a p sublevel with one orbital from an s sublevel. One p orbital remains unchanged.

Intermolecular Forces the higher the force of attraction between molecules, the harder it is to pull them apart. high BP = large attractive forces low BP = small attractive forces the strongest intermolecular forces act between polar molecules. Dipole: formed due to differences in electronegativity. A dipole is represented by an arrow pointing towards the negative pole (more electronegative element).

dipole-dipole forces: the negative region of one molecule is attracted to the positive region of another molecule. Balloon- Water Explanation a polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons  “induced dipole”. Induced Dipole

Hydrogen Bonding a strong type of dipole-dipole force that accounts for the unusually high BP of hydrogen containing compounds. O H London Dispersion Forces a weak intermolecular attraction resulting from instantaneous dipoles. all molecules experience London Forces, however they are the only force of attraction among noble-gases, nonpolar, and slightly polar molecules. London Forces