Chapter 12 Chemical bonding

12.1 Types of Chemical Bonds Bonds: a force that holds groups of two or more atoms together and makes them function as a unit Required 2 e- to make a bond Bond energy: amount of energy required to form or to break the bond

Ionic Bonding Occurs in ionic compound Results from transferring electron Created a strong attraction among the closely pack compound

Covalent Bonding Formation of a covalent Bond Two atoms come close together, and electrostatic interactions begin to develop Two nuclei repel each other; electrons repel each other Each nucleus attracts to electrons; electrons attract both nuclei Attractive forces > repulsive forces; then covalent bond is formed

12.2 Electronegativity Electronegativity (EN): the ability of an atom in a molecule to attract the shared electron in a bond Metallic elements – low electronegativities Halogens and other elements in upper right-hand corner of periodic table – high electronegativity

Polarity Polar covalent bonds – the bonding electrons are attracted somewhat more strongly by one atom in a bond Electrons are not completely transferred More electronegative atom: δ- . (δ represents the partial negative charge formed) Less electronegative atom: δ+

Relationship Between Electronegativity and Bond Type Predicting bond polarity Atoms with similar electronegativity (Δ EN <0.4) –form nonpolar covalent bond Atoms whose electronegativity differ by less than two (Δ EN < 2) – form polar covalent bonds Atoms whose electronegativity differ by more than two (Δ EN > 2) – form ionic bonds

Examples Using the electronegativity values given in Figure 12.3, arrange the following bonds in order of increasing polarity: H-H, O-H, Cl-H and F-H For each of the following pairs of bonds, choose the bond that will be more polar a. H-P, H-C b. N-O, S-O c. O-F, O-I d. N-H, Si-H

12.3 Polarity and Dipole Moment a vector quantity from the center of the positive charge to the center of negative charge Represents with an arrow E.g Draw the dipole moment for HF, H2O, HCl, OF

13.4 Stable Electron Configurations and Charges on Ions Atoms in stable compounds almost always have a noble gas electron configuration Predicting Formulas of Ionic Compound Electrons lost by a metal come from the highest-energy occupied orbital Electrons gained by a nonmetal go into lowest-energy unoccupied orbital

Examples Predicting formulas of Ionic compound by showing how they loses or gains electrons Ca and O Li and S

12.5 Ionic Compound Lattice energy (U) – the sum of the electrostatic interaction energies between ions in a solid Refer to the breakup of a crystal into individual ions

12.6 Lewis Structures represents how an atom’s valence electrons are distributed in a molecule Show the bonding involves (the maximum bonds can be made) Try to achieve the noble gas configuration

Rules Duet Rule: sharing of 2 electrons E.g H2 H : H Octet Rule: sharing of 8 electrons Carbon, oxygen, nitrogen and fluorine always obey this rule in a stable molecule E.g F2, O2 Bonding pair: two of which are shared with other atoms Lone pair or nonbonding pair: those that are not used for bonding

Lewis Structures B C N O F use common bonding patterns C = 4 bonds & 0 lone pairs, N = 3 bonds & 1 lone pair, O= 2 bonds & 2 lone pairs, H and halogen = 1 bond, Be = 2 bonds & 0 lone pairs, B = 3 bonds & 0 lone pairs often Lewis structures with line bonds have the lone pairs left off their presence is assumed from common bonding patterns structures which result in bonding patterns different from common have formal charges B C N O F Tro, Chemistry: A Molecular Approach

Rules Create a skeletal structure using the following rules: Hydrogen atoms (if present) are always on the “outside” of the structure. They form only one bond The central atom is usually least electronegative. It is also often unique (i.e,. the only one atom of the element in the molecule). Remember, there might be no “central” atom. Connect bonded atoms by line (2-electron, covalent bonds

Rules Place lone pairs around outer atoms (except hydrogen) so that each atom has an octet Calculate the number of electrons you haven’t used. Subtract the number of electrons used so far, including electrons in lone pair and bonding pairs, from the total in Step 1. Assign any remaining electrons to the central atom as lone pair

Rules 5. If the central atom is B (boron) or Be (beryllium), skip this step If the central atom has an octet after step 4, skip this step If the central atom has only 6 electrons, move a lone pair from an outer atom to form a double bond between outer atom and the central atom If the central atom has only 4 electrons, do Step 5a to two different outer atoms (i.e, form two double bonds) or twice to one outer atom (i.e., form one triple bond)

Examples Draws a dot Lewis structure for Na, O, Cl, H2O, HCl, CCl4

12.7 Lewis Structures of Molecules with Multiple Bonds Recall: Elements typically obey the octet rule; they are surrounded by eight electrons single bond: involves two atoms sharing one electron Double bond: involves two atoms sharing two pair of electrons Use 6N + 2 Rule N = number of atoms other than Hydrogen

Molecular Structure: The VSEPR Model VSEPR: Valence shell electron pair repulsion Predicting the molecular structure of molecules formed from nonmetals Bonding and nonbonding pairs (lone pairs) around given atom are positioned as far as possible

Examples Give the Lewis structure for the following CF4, CO, SO42- NH4+

Vector Addition Tro, Chemistry: A Molecular Approach

Tro, Chemistry: A Molecular Approach