Chem 121 1st Edition Lecture 13 Outline of Last LectureI. Nomenclature- Anions / Naming Simple CompoundsII. Formula for Ionic Compounds Outline of Current Lecture I. Chemical BondingII. Bond LengthIII. Bond EnergyIV. Bond OrderV. Electronegativity (EN)VI. Bond Polarity and Dipole MomentVII. Lewis StructuresCurrent LectureChemical Bonding- Atoms get close to share valence e - to lower energy level- Ionic Bonding: The electrostatic attraction between oppositely charged ionso Metal: Cation (Loses electrons)o Nonmetal: Anion (Accepts electrons)- Covalent Bonding: electrons forming pairs are shared by atomso Nonmetal- Nonmetal- Covalent Bonds: valence electron pair shared between atomso Valence electron pairs can either share equally or unequallyThese notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.o If shared unequally, one electron pair is spending more time in one atom- Ionic Bonds: Valence electron pair owned by one atom, where atoms become ionso Ionic bonds are extreme cases of covalent bondso The type of bond depends on how strong the attraction is between the bonding pairsBond Length- Distance between the nuclei of two bonding atoms- Sum of radii of bonding atoms or bonding ions- Bond length and bond strength are reverses of each otherBond Energy (BE, or bond strength)- Energy required to break a bond into gaseous atomso A-B (g)  A (g) + B (g) BEA-B > 0  gain- Forming a bond- Release energyo A (g) + B (g)  A-B (g) -BEA-B < 0  lose- Bond order and Bond length are the same- Type of bondso Single bond: one bonding pair of electrons (bond order: 1)o Double bond: two bonding pairs of electrons (bond order: 2)o Triple bond: three bonding pairs of electrons (bond order: 3)Bond Order- Number of electron pairs being shared by bonding atoms- Relation between bond order, bond energy, and bond length: “ For a given pair of atoms: The higher the bond order, the shorter the bond length, and the higher the bond energy”  SUPER IMPORTANT!Electronegativity (EN)- Definition: “ The ability of an element to attract a sharing or bonding electron pair to itself within a covalent bond.”o EN  [ Ei1 + (-EA)]- Period trend- Results of Zo Down a group:  EA gets weaker (less negative) Ei1 decreases EN decreaseso Cross a period (L to R) EA gets stronger (more negative) Ei1 increases EN increaseso IN GENERAL: ENmetal < ENnonmetal Electron “hogs” (greatest values)- F ENF: 4.0- O ENO: 3.5- N ENN: 3.0- Cl ENCl: 3.0 H (Least EN nonmetal) ENH: 2.1Bond Polarity and Dipole Moment- Polarity of Bonds- Pure Covalent Bonds (nonpolar covalent bonds): Bonds between two nonmetal atoms with equal EN o Bonding pair(s) are equally sharedo NO polarity in bonds, no partial chargeo Diatomic: Br2, I2, N2, Cl2, H2, O2, F2o Multiatomic: O3, NCl3, PH3, CS2o EN A-B = 0 (C-H bond is considered as a pure covalent)- Polarity Covalent Bondso Bonding Atoms with different EN (EN doesn’t equal 0)o Bonding pair unequally shared – more electron density on more EN atom: less electron density on less EN atomo Partially charged centers on atoms Partial negatively charged center  more EN atom Partial positively charged center  less EN atomo 0 < EN A-B < 2.0o “ The greater EN, the more polar the covalent bond, the larger the partial charges, the more ionic character of the covalent bond.” - Ionic Bonds- between metal and nonmetal atomso Bonding pair owned by one atom (an extreme case of covalent bond)o EN A-B > 2.0o Complete (full) charges on ions Comparison between different bond types::o Pure covalent bonds have no charge on the atomo Polar covalent bonds have partial charge on the atomo Ionic bonds have whole (all) charge on the atomLewis Structures Procedures- (Key: take valence electrons from all bonding atoms, then redistribute)1. Add up valence electrons from each atom2. Set atomsa. Central Atom: Less EN, except H Subscript: number of central atomsb. Surrounding (terminal) atoms: rest of the atom, symmetrically arranged (Acid: H attaches to O)3. Distribute electronsa. Single bond between each pair of atomsb. Rest of electrons:i. First—surrounding atoms: duet for H; Octet for rest of atomsii. Second—Central atom4. Octet for Central Atom:- If no octet, form double or triple bonds with terminal atoms- Move 1 or more lone electron pairs from terminal atoms to bonding position between