Chemical Bonding. Drawing Electron Dot diagrams/lewis Structures for simple particles: Lewis diagrams show only valence electrons:

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1 Chemical Bonding Chemical Bonding results from the attraction between the atoms or ions of a substance. Atoms bond in order to reduce their individual energy and gain stability. They gain this stability by gaining, losing or sharing electrons to achieve a stable octet. Types of Bonding 1. Ionic Bonding = the electrostatic attraction between positive and negative ions which are formed when a metal transfers its electrons to a non-metal 2. Covalent Bonding = the sharing of pair(s) of electrons between two or more non-metals. Many atoms do so to complete the s and p subshells in their valence shell. The octet rule states that when atoms form covalent bonds, they tend to share sufficient electrons to achieve a valence shell with eight electrons Lewis Theory of Bonding (4.1) Key Ideas of the Lewis theory of Bonding Atoms & ions are stable if they have a noble gas configuration of electrons Electrons are more stable if they are paired Atoms form chemical bonds to achieve stability A stable octet may be achieved by exchanging of electrons between metals and non-metals or the sharing of electrons between non-metals The sharing of electrons results is a covalent bond Drawing Electron Dot diagrams/lewis Structures for simple particles: Lewis diagrams show only valence electrons: EXAMPLES: Drawing Lewis Structures for more complex particles (see rules) EXAMPLES: Exceptions to Lewis Octet Rule (see sheet Lewis Structures & the Octet Rule) Comparing Lewis Structures with Quantum Mechanics (see Table 1, pg. 226, Nelson 12)

2 Covalent Bonding and the Valence Bonding Theory (4.2) Lewis Structures are useful for electron bookkeeping, but they tell us nothing about why covalent bonds form or how the electrons manage to be shared between atoms. According to the valence bond theory, which was developed by Linus Pauling, a bond between two atoms is formed when a pair of electrons is shared by two overlapping orbitals; one orbital from each atom is joined in the bond. The shared electron pair becomes concentrated in the region of overlap and helps glue the nuclei together. Therefore atoms tend to position themselves so that maximum amount of orbital overlap occurs. Summary of Valence Bond theory: A half filled orbital in one atom can overlap with a half filled orbital of a second orbital to form a new, bonding orbital The new, bonding orbital contains a pair of electrons of opposite spin The total number of electrons in a bonding orbital must be two When atoms bond, they arrange themselves in space to achieve the maximum overlap of their half-filled orbitals. Maximum overlap produces a bonding orbital with the lowest energy (i.e. greater stability) Examples:

3 Hybridization = formation of Hybrid Orbitals Hybridized orbitals are the result of the mixing of the simple s, p, d, f orbitals to form more stable atomic orbitals. When these orbitals mix they are able to create a new set of orbitals that are identical in energy, shape and size. This mixing of atomic orbitals produces better overlap of the new hybrid orbitals, which result in the formation of stronger bonds. (See Sheets on hybridization). Sigma ( ) & Pi Bonds ( ) Sigma bonds are formed by the end-to-end overlap atomic orbitals; sigma bonds can be thought of as single bonds; they permit free rotation around the bond axis. e.g. Pi Bonds are formed by the side-to-side overlap of p orbitals Pi bonds are the second and third lines in the structural diagram of double and triple covalent bonds e.g hydrocarbons: C 2 H 4 results from partial hybridization of the atomic orbitals in the carbon atoms to give three sp 2 hybrid orbitals which forms sigma bonds and one unhybridized p orbital which participates in a pi bond C 2 H 2 results form the partial hybridization of the atomic orbitals in the carbon atoms to give TWO sp hybrid orbitals which forms sigma bonds and TWO unhybridized p orbital which both participate in a pi bond

4 VSEPR Theory (4.3) Understanding the structure (including shape) a molecule is vital to explaining its chemical and physical properties. The shape of a molecule is investigated through x-ray analysis of crystals (i.e. X-ray crystallography). Lewis structures account for the bonding of each atom in a molecule, but not the 3-D shape of the unit. The VSEPR (Valence Shell Electrons Pair Repulsion Theory) was created by Gillespi and Nyholm, proposes that the arrangement of atoms around a central atom in a molecule depend of the repulsion between all the electron pairs in the valence shell of the central atom. These electron pairs (in their orbital) stay as far apart as possible (have the greatest angles of separation) so as to minimize the repulsion of their negative charge. Summary of VSEPR Theory Only the valence shell of the central atom (i.e the atom or atoms in a molecule that has/have the most bonding electrons (i.e. can form the most bonds) are important for molecular shape Valence shell electrons are paired or will be paired in a molecule or polyatomic ion Bonded pairs of electrons and lone pairs are treated approx. equally* Valence shell electrons pairs repel each other electrostatically The molecular shape is determined by the positions of the electron pairs when they are a maximum distance (greatest angle) apart * Orbitals with lone pairs on the central atom occupy more space than those with bonding pairs. This is because lone pairs are under the influence of only one nucleus, so they are bigger and puffier than the shared electron pairs which are attracted to two nuclei (and thus are more elongated) Summary on how to determine shapes of molecules Step 1: Draw the Lewis Structure for the molecule, including the electron pairs around the central atom Step 2: Count the number of bonding pairs (bonded atoms) and lone pairs around the central atom Step 3: Refer to Molecular Shape sheet and use the number of paired electrons to predict the shape of the molecule. N.B. Treat double or triple bonds as single bonds. Examples:

5 Review of Trends in the periodic Table: Electronegativity & Covalent Bonding (4.4) To examine covalent molecules, we must understand the quantitative measure of electronegativity. Electronegativity is the measure of electron attracting ability of the atoms in a molecule. The electronegativity value of an atom increases across a period and decreases across a group. Non-Polar Covalent Bonds = form when there is equal sharing of electron pairs between 2 nuclei. Happens when the electronegativity difference between bonding atoms in zero or close to zero (<0.8) Polar Covalent Bond = form when there is an unequal sharing of electron pairs between 2 nuclei. This occurs when the electronegativity difference is greater than 0.8 but less than 1.7. Due to this electronegativity difference, a dipole is established within the molecule Ionic Bonds = form if the electronegativity difference between the two atoms is greater than 1.7 Bonding Continium (see pg. 252 in Nelson 12) Polar Molecules vs Non-polar molecules Non-Polar molecule = one with only non-polar bonds or polar bonds whose dipole cancel each other Polar Molecules = molecule with polar bonds that do not cancel Both the shape and the polarity of the bonds are needed to determine if a molecule is polar or not. In all symmetrical molecules, the dipoles cancel, so the molecule is non-polar Intermolecular Forces & how they explain Physical Properties (eg. Boiling points, solubility) (4.5) The Structure & Properties of Solids (4.6): Properties are due to the kinds of forces/bonds in these solids a) Ionic Crystals b) Metallic Crystals c) Molecular Crystals d) Covalent Network Crystals