Chapter 16 Covalent Bonding

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1 Chapter 16 Covalent Bonding Reading Assignment C16! 1. Read C16 pp and while reading, continue building your science vocabulary table that includes all terms in bold face type and all terms you are unfamiliar with or unsure of. Your vocabulary table needs to include three things: 1. The term 2. The definition 3. A picture of what the term means to you. 2. On page 469, organize the Concept Map into a form that correctly connects the major ideas of the chapter. Lewis Theory of Covalent Bonding- The driving force of bond formation is the desire of each atom in a molecule to attain an octet of electrons in its valence shell. I. Covalent Bond - A chemical bond that is formed from a pair of shared electrons between two atoms. a. Caused by a small difference in electronegativities of the bonded atoms. i. General rules defining covalent bonds: 1. If Δ electronegativities < 2.0 (1.7). 2. If Allen Factor is < 0.4. ii. General characteristics of covalent bonds: 1. Typically between 2 nonmetals. 2. Occur from an overlapping of orbitals where shared electrons are attracted by both nuclei. 3. Very strong bonds. iii. Lewis Structures- Diagrams that show how the valence electrons are arranged among the atoms in the molecule showing how the atoms bond. 1. The valence electrons about an atom in a molecule as either: a. Shared electrons (bonded electrons). b. Unshared electrons (lone pair electrons). 2. Guidelines for writing Lewis Structures. a. Draw a simple structural formula using single covalent bonds. The least electronegative atom serves as the central atom. ** Never use Hydrogen as the central atom** b. Determine the total number of valence electrons in the molecule or ion. i. For molecules- sum of the valence electrons on each atom. ii. For cations- sum of the valence electrons on each atom the charge of the cation. iii. For anions- sum of the valence electrons on each atom + the charge of the anion.

2 II. c. Deduct 2 valence electrons for each bond from 2., then distribute the remaining about each atom as lone pairs to fill the octet (8). If too few electrons exist, convert single bonds to multiple bonds. i. Most multiple bonds occur between Carbon(s), Oxygen(s), Nitrogen(s), Sulfur, & Phosphorus. iv. Formal Charge - The hypothetical charge assigned to each atom in a Lewis structure. 1. Rules for calculating Formal Charge. a. One half of the electrons in a bond are assigned to each atom. (1 each for single, 2 each for double, etc.). b. Both electrons of an unshared pair are assigned to the atom to which the unshared electron belongs. 2. Formal Charge = the number of valence electrons in the free atomic state minus the number of assigned electrons of the atom in a molecule. a. The sum of the formal charges of a molecule should equal zero, while the sum of the formal charges of an ion should equal the charge of the ion. b. Formal charges are used to predict structural formulas for varying Lewis structures for any molecule. 3. Rules for predicting Lewis Structures using Formal Charges. a. A Lewis structure in which all formal charges in a molecule are zero is preferable to one in which some formal charges are not zero. b. If Lewis structures must have nonzero formal charges, the one with the lowest number of nonzero formal charges is preferable, and a Lewis structure with no more than one large formal charge (- 2,+2,etc.) is preferable to one with several c. Lewis structures should have adjacent formal charges of zero or have opposite sign. d. When we must choose among several Lewis structures with similar distributions of formal charges, the structure with the negative formal charges on the more electronegative atoms is preferable. Molecular Formulas- Identifies the elements and bonds in a molecule. a. Writing & Naming molecular formulas. i. Write the least electronegative atom first. ii. Use prefixes to designate the ratio of atoms in a molecule. See table below.

3 iii. Change the second name ending to (-ide). 1- mono 6- hexa 2- di 7- hepta 3- tri 8- octa 4- tetra 9- nona 5- penta 10- deca b. Diatomic molecules Those atoms that when in elemental state exist as diatomic molecules (2 bonded atoms). i. The "super seven": H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 c. Types of Covalent Bonds. i. Single covalent- sharing of 1 pair of electrons between bonding nuclei. ii. Double covalent- sharing of 2 pairs of electrons between bonding nuclei. iii. Triple covalent- sharing of 3 pairs of electrons between bonding nuclei. d. Coordinate Covalent Bonds - A covalent bond where 1 atom provides both bonding electrons- Can be drawn as an arrow. e. 5. Resonance - A situation where two or more Lewis structures for a molecule can be written which only differ in the arrangement of electrons (shared and unshared). f. Exceptions to the Octet Rule for Covalent Molecules i. Magnetic charges - due to unpaired electrons. Remember, electron spin creates a magnetic field. 1. Diamagnetic Compounds- compounds that have no unpaired electrons. a. These compounds are not affected by a magnetic field. 2. Paramagnetic Compounds- compounds that have unpaired electrons therefore are affected by magnetic fields. ii. Octet Expansion & Contraction. 1. Elements in 3 rd or higher energy level can obtain more than 8 valence electrons in the valence shell of the central atom, due to d-orbital hybridization. 2. Boron and Aluminum can exhibit configurations with less than 8 valence electrons. g. Correlating Oxidation Numbers in Ionic and Covalent Bonds. i. The oxidation number of any atom in its elemental form is Zero. ii. The oxidation number of a monatomic ion is considered equal to the charge on the ion. iii. The oxidation number of Fluorine is always 1. iv. Alkali Metals have an oxidation number of +1. v. Alkaline Earth Metals have an oxidation number of +2.

4 vi. Halogens have oxidation numbers of 1 when paired with less electronegative atoms. vii. Oxygen s oxidation number is usually 2 except: 1. When bonded with Fluorine, the oxidation number is positive. Ex. OF 2, oxygen is O In peroxides (O 2-2 ) the oxidation number is 1 3. In superoxides (O 2-1 ), the oxidation number is ½ Ex. KO 2. viii. Hydrogen is assigned a 1 with less electronegative atoms and +1 with greater electronegative atoms. ix. The sum of the oxidation numbers of all the atoms in a compound is zero. The sum of the oxidation numbers in an ion is equal to the charge of the ion. h. Valence Shell Electron Pair Repulsion Theory VSEPR i. Used to predict geometric shape of a molecule due to the repulsion of electrons (bonded and lone pairs). ii. Valence Bond Theory- Describes covalent bonds as the overlapping of orbitals which produces a mutual attraction on bonded electrons by adjacent nuclei. 1. Sigma Bond (σ)- covalent bond where electron density is concentrated in the region along the inter-nuclear axis. a. Between s/s orbitals, s/p orbitals, and directional p/p bonds. 2. Pi Bond (π)- covalent bond resulting from the side-by-side bonding of p orbitals. The amount of overlap is less than a sigma bond; therefore the pi bond is weaker. 3. Multiple bonds: a. single- sigma b. double- sigma & 1 pi c. triple- sigma & 2 pi bonds iii. Hybridization- Hybrid orbitals are created from the mixing of atomic orbitals that are covalently bonded. 1. Facts about hybridization: a. Hybrid orbitals are formed only in covalently bonded atoms. b. Hybrid orbitals have different shapes & orientations than the atomic orbitals. c. All orbitals in a set of hybrid orbitals are equivalent and form identical bonds. d. The number of Hybrid orbitals is equal to the number of atomic valence orbitals that are used to form the hybrid orbitals. e. The type of Hybrid orbitals on a bonded atom depends on the geometry of its unshared electron pairs and atoms bonded to it.

5 f. Sigma bonds occur from the overlapping of hybridized orbitals and pi bonds occur from the overlap of unhybridized orbitals. 2. Types of hybridization: a. sp 3-1 s and 3 p orbitals producing 4 hybrid orbitals. i. Produces a tetrahedral arrangement. b. sp 2-1 s and 2 p orbitals producing 3 trigonal planar hybridized orbitals and 1 p is left unhybridized. c. sp- 1 s and 2 p orbitals producing 2 linear hybridized orbitals and 2 unhybridized orbitals. d. others- for those of the octet exceptions, due to hybridization of n=3 (or higher) d-type orbitals. i. sp 3 d- (1 s, 3 p, & 1 d) producing trigonal bipyramidal shape. ii. sp 3 d 2 - (1 s, 3 p, & 2 d) producing octahedral shape. 3. Molecular Structure 3-dimensional arrangement of atoms. a. Bond distance (radii)- the distance between 2 bonded nuclei- measured in pm (10-12 m) or (10-10 m). b. Bond angle- the angle produced between 2 bonded that share a common atom. iv. Using VSEPR for predicting shapes- To predict the geometry you must 1 st draw proper Lewis structures 1. AX m (LP) n : a. A - represents the central atom. b. X - peripheral (bonded) atoms. m is the number of peripherally bonded atoms. c. LP - Lone Pairs about the central atom. n is the number of lone pairs. i. Polar Covalent Bonds - Covalent bonds that result in partially charged positive ends and partially charged negative ends due to differences in electronegativities. i. Nonpolar Covalent Bond - No net partial charges due to similar electronegativities between bonded atoms. 1. Electronegativity - a measure of the attraction on an electron by an atom in a chemical bond. a. Δ electronegativity = e.n. 1 e.n. 2 gives a rough estimate of polarity of a bond.

6 Electronegativity Differences and Bond Types Δ e.n Bond Type Nonpolar covalent Polar covalent a Strongly polar covalent > 2.0 a Ionic a -If the compound is made from a metal/nonmetal and Δ e.n > 1.6 then it's ionic ii. Dipolarity - Polarity in molecules (electronic lopsidedness) is dependent upon bonding and geometry. 1. Dipole moment - (DeBye dipole moment µ D ) found by placing a molecule into an electric field. a. A polar molecule will align itself to the polarity of the field. i. µ D = 0 for nonpolar molecules ii. µ D > 0 for polar molecules Examples: Identifying Dipole moments Molecule Bond Polarity of molecule CO 2 Polar Nonpolar NH 3 Polar Polar H 2 O Polar Polar O 3 Nonpolar Polar 2. Bond dissociation energy - the energy required to break a single bond between two atoms. (typically calculated when molecule is in gaseous state). a. Described as the enthalpy change of the endothermic reaction. b. Ex. XY (g) X (g) + Y (g) D X-Y = ΔH (D X-Y is the bond energy) i. Kilojoule = 1000 joules, 1 cal = joules. c. **The strength of a bond between two atoms increases as the number of electron pairs in the bond increases.** 3. Comparison of Covalent vs. Ionic Compounds: Covalent Ionic a. Are composed of discrete a. Are composed of (+) and (-) ions

7 particles called molecules. The molecules are made up of atoms held together by covalent bonds (shared electrons) b. May exist as solids, liquids, or gases c. Electronegativity differences are small stacked aroundeach other. The ions are held together by ionicbonds (electrostatic attractions). b. Exist as brittle, crystalline solids at room temps. c. Electronegativity differences are high d. Composed of 2 non-metals d. Binary compounds are formed from the reaction of a metal and a non-metal. e. Contain directional bonds and have definite bond angles f. Usually have lower melting and boiling points Ternary and higher compounds are formed from polyatomic ions e.contain no directional bonds. Ionic compounds stacks of ions f. Exhibit high melting and boiling points

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