Chapter 12 Covalent Bonding

1. Chapter 12 Covalent Bonding • Many substances in nature are molecular. • Molecules are composed of atoms combined together. • Unlike ionic compounds, which are formed by oppositely attracted charged atoms, molecules are formed by neutral atoms. • These atoms combine by sharing outer electron orbitals to achieve noble gas electron configuration.

2. Covalent Bonding The simplest molecules are the diatomic elemental substances H2, N2, O2, F2, Br2, I2, Cl2. In the hydrogen molecule, both hydrogen atoms provide a single electron into the s-orbital to achieve a similar electron configuration as helium. This is referred to as a single covalent bond because one pair of electrons are shared.

3. Covalent Bonding • Oxygen has 6 electrons in the valence shell (outer s and p orbitals). • 2 electrons are needed to become isoelectronic with Neon. • Oxygen supplies the two electrons with its partner supplying two. This is a double covalent bond.

4. Covalent Bonding • Water is a covalently bonded molecule made up of 2 hydrogen and 1 oxygen atoms. • The hydrogen atoms provide 1 electron each to the oxygen as well as oxygen sharing an electron to each of the hydrogens.

5. Unequal Electron Sharing • Electronegativity • The element’s attractiveness to an electron pair in a covalent bond • The values range from 4.0 (Fluorine) to .70 (Francium) • The larger the difference between elements in a covalent bond, the more the electrons are moved closer to the more electronegative element.

6. Electronegativity

7. Bond Polarity • A type of covalent bond between two atoms in which electrons are shared unequally. Because of this, one end of the molecule has a slightly negative charge and the other a slightly positive charge.

8. Lewis Dot Structure Notation Lewis Dot structures are a shorthand notation of keeping track of valence electrons available for sharing in covalent bonds. Each dot adjacent to the element symbol represents valence electrons.

9. Lewis Dot Structures