1	Some Notes on Covalent Bonding and Lewis Structures
2	Covalent bonds: formed from sharing of two electrons, usually one donated from each of the two bonding atoms. Sharing electrons is one way that atoms can satisfy the "Octet Rule" which as stated by Gilbert Lewis "atoms, by sharing electrons to form an electron-pair bond, can acquire a stable, noble-gas structure".
3	Covalent bonds: There are 3 kinds of covalent bonds based on number of pairs of electrons shared between the two atoms: single covalent bond - 1 shared pair double covalent bond - 2 shared pairs triple covalent bond - 3 shared pairs
4	Covalent bonds: Comparison of different kinds of bonds Bond energy: single < double < triple Bond length: single > double > triple
5	Lewis Structures: Simplest covalent bond is between two hydrogen atoms. Electron-dot (or Lewis) symbol for hydrogen is H^. The two electrons from two hydrogen atoms pair up to make H:H. The pair of dots can be replaced by a dash to represent a bond, as in H-H Hydrogen as the element exists as these "diatomic molecules" and its molecular formula is shown as H₂.
6	Lewis Structures: halogens also exist as diatomic molecules. Fluorine, chlorine, bromine and iodine all have the same type of Lewis structure seven outer electrons (2 on the right side, 2 on the top, 2 on the left side, and 1 on the bottom, so two F atoms share their unpaired electrons with each other to form the single covalent bond, F:F or F-F, to get F₂. The rest of the pairs of electrons around each halogen are called "unshared pairs". The number of paired and unpaired electrons around an atom play an important part in determining the shape of a molecule, which in turn contribute to chemical and physical properties of the molecule.
7	Lewis Structures: Oxygen atoms have six outer level electrons each ( 2 on the right side, 2 on the top, and 1 each on left side and bottom) share two pairs of electrons and thus have a double bond between them This is shown by O::O or O=O for O₂.
8	Lewis Structures: Nitrogen atoms have only five outer electrons have 2 on the right side, and 1 each on the top, left side and bottom. Two nitrogen atoms share three electrons each, to form three pairs between them. This bonding is written as :N:::N:, N=N for N₂
9	Lewis Structures: Things get a bit more complicated when atoms of different kinds of elements bond. One of the simplest bonds is between hydrogen and chlorine. The single unpaired electron on H pairs up with the unpaired electron on the chlorine result is H:Cl or H-Cl, or HCl.
10	Writing Lewis Structures for simple molecules. 1. Count the number of valence electrons in each of the atoms in the molecule. 2. Count the total number of electrons that are needed to give each atom in the molecule an octet in the valence level, or two electrons for hydrogen.
11	Writing Lewis Structures for simple molecules.
12	Writing Lewis Structures for simple molecules. 4. Divide the number of shared electrons by 2 to give the number of electron pairs, thus the number of bonds between all of the atoms in the molecules. A double or triple bond count as two or three bonds because of the number of electron pairs they represent.
13	Writing Lewis Structures for simple molecules. CH₂O C₂H₄ H₂O CO₂ ClO₃^1-
14	Rules to remember when writing Lewis structures: H is always terminal, only one single bond per hydrogen Oxygen almost always has two shared pairs and two unshared pairs of electrons. The two shared pairs may be in two separate single bonds or in one double bond
15	Rules to remember when writing Lewis structures: Carbon atoms will always have 4 bonds, either as 4 separate single bonds, or a combination of single, double and/or triple bonds so that there are 4 pairs of electrons around each carbon atom in the molecule. Carbon atoms may single, double, or triple bond to each other, single or double bond to oxygen atoms, or single, double, or triple bond to nitrogen atoms.
16	Rules to remember when writing Lewis structures: Halogen atoms form single bonds only, and usually form just one bond. Example of an exception is the chlorate, ClO₃^1-, ion Chlorine is the central atom with the 3 oxygen atoms bonded to it
17	Resonance and Resonance Structures Sometimes 2 or more structures can be drawn for a molecule…all of which are correct When this difference is the placement of double or triple bonds, how does one decide which one is correct? Concept of resonance is used. Our models have limitations.
18	Resonance and Resonance Structures SO₃^2- SO₃ Examples SO₂ NO₃^1-
19	Resonance and Resonance Structures Examples C₆H₆
20	Resonance and Resonance Structures Notice that in C₆H₆and C₂H₄ there is no single central atom Consider each carbon atom individually A molecule can have more than one kind of geometry around different atoms. C₂H₄O₂
21	Electronegativity and Polar Bonds
22	What is electronegativity? Electronegativity is a measure of the attraction of an atom for electrons in a covalent bond. The most commonly used electronegativity scale is Pauling's. It is the one we will usually use.
23	What is electronegativity? Fluorine, the most reactive non-metal, is assigned the highest value since it has the greatest attraction for the electron being shared by the other element. Oxygen is also highly electronegative and has a strong attraction for electrons. Metals have low electronegativities since they have weak attraction for any shared electrons.
24	What is electronegativity? When two unlike atoms are convalently bonded, the shared electrons will be more strongly attracted to the atom of greater electronegativity. Such a bond is said to be polar. A polar bond results in the unequal sharing of the electrons in the bond.
25	What is electronegativity? The presence or absence of polar bonds within molecule plays a very important part in determining chemical and physical properities of those molecules. Some of these properties are melting points, boiling points, viscosity and solubilities in solvents
26	Predicting Bond Types The difference in electronegativities of two elements can be used to predict the nature of the chemical bond. Bond type can be described as belonging to one of three classes: nonpolar covalent polar covalent ionic
27	Predicting Bond Types When differences are 1.7 or greater, the bond is usually ionic. Less than 1.7, the bond is usually covalent, and unless the difference is less < 0.4 the bond has some degree of polarity. Differences of < than 0.4 are considered to be nonpolar.
28	Predicting Bond Types Note the electronegativities of elements from left to right across periods and down groups. What is the trend for each of these? Notice that Pauling didn't assign a value to the Noble gases
29	Predicting Bond Types Across a period: the electronegativities generally increase from left to right across a period with the Group VII element having the highest value for the period. Down a group: the electronegativities generally decrease from top to bottom down a group. Francium is the element with the lowest electronegativity.
30	Predicting Bond Types C-Cl 0.61 polar P-Cl 0.97 polar N-O 0.40 polar C-S 0.03 non-polar C-H 0.35 non-polar C-O 0.89 polar O-H 1.24 polar N-H .84 polar
31	Molecular Shapes
32	Molecular Shapes
33	Molecular Shapes
34	Molecular Shapes
35	Molecular Shapes
36	Molecular Shapes
37	Molecular Shapes The polarity of a molecule depends on the presence of polar bonds and the shape of the molecule. Asymmetrically shaped molecules with polar bonds are polar Symmetrically shaped molecules are not polar if the polar bonds are the same
38	Molecular Shapes
39	Molecular Shapes