Chemistry B2A Chapter 12 Chemical Bonding

Chemical Bonds Ionic bonds 2. Covalent bonds 3. Metallic bonds 4. Hydrogen bonds 5. Van der Waals forces

Chemical Bonds Ionic bonds 2. Covalent bonds

Review Level 1 Maximum 2 electrons in valence level Hydrogen and Helium main-group elements 1A – 8A Other levels Maximum 8 electrons in valence level

Na: 1s2 2s2 2p6 3s1 Na+: 1s2 2s2 2p6 Ne: 1s2 2s2 2p6 Octet & Duet rules Noble gases (Stable) Goal of atoms Filled valence level Na+: 1s2 2s2 2p6 + e- Na: 1s2 2s2 2p6 3s1 Ne: 1s2 2s2 2p6 Ar: 1s2 2s2 2p6 3s2 3p6

Na: 1s2 2s2 2p6 3s1 Na+: 1s2 2s2 2p6 Ne: 1s2 2s2 2p6 Octet & Duet rules Noble gases (Stable) Goal of atoms Filled valence level Na+: 1s2 2s2 2p6 + e- Na: 1s2 2s2 2p6 3s1 Ne: 1s2 2s2 2p6 + e- Cl: 1s2 2s2 2p6 3s2 3p5 Cl-: 1s2 2s2 2p6 3s2 3p6 Ar: 1s2 2s2 2p6 3s2 3p6

Mg: 1s2 2s2 2p6 3s2 Mg2+: 1s2 2s2 2p6 Ne: 1s2 2s2 2p6 Octet & Duet rules Noble gases (Stable) Goal of atoms Filled valence level + 2e- Mg: 1s2 2s2 2p6 3s2 Mg2+: 1s2 2s2 2p6 Ne: 1s2 2s2 2p6

Mg: 1s2 2s2 2p6 3s2 Mg2+: 1s2 2s2 2p6 Ne: 1s2 2s2 2p6 O: 1s2 2s2 2p4 Octet & Duet rules Noble gases (Stable) Goal of atoms Filled valence level + 2e- Mg: 1s2 2s2 2p6 3s2 Mg2+: 1s2 2s2 2p6 Ne: 1s2 2s2 2p6 + 2e- O: 1s2 2s2 2p4 O2-: 1s2 2s2 2p6 Ne: 1s2 2s2 2p6

Cation (Y+): Na+ Li+ Ca2+ Al3+ Metals: lose 1, 2 or 3 e- Cation (Y+) Ions Nonmetals: gain 1, 2 or 3 e- Anion (X-) Number of protons and neutrons in the nucleus remains unchanged. Cation (Y+): Na+ Li+ Ca2+ Al3+ Anion (X-): Cl- F- O-2

Transition elements 1A 2A 3A 4A 5A 6A 7A 8A

Two problems of Octet rule Octet rule cannot be used for type II cations (most of the transition and inner transition elements). Fe2+ Fe3+ Cu1+ Cu2+ 2. Ions of period 1 and 2 elements with charges greater than +2 are unstable. C C4+ C4- B B3+ unstable unstable

Ionic bonds Metal-Nonmetal Na: 1s2 2s2 2p6 3s1 Cl: 1s2 2s2 2p6 3s2 3p5 Anion Cation Na+: 1s2 2s2 2p6 Cl-: 1s2 2s2 2p6 3s2 3p6

Sodium (Na) NaCl Chlorine (Cl)

matter are neutral (uncharged): total number of positive charges = total number of negative charges Na+ Cl- NaCl Ca2+ Cl- CaCl2 Al3+ S2- Al2S3 Ba2+ O2- Ba2O2 BaO

Naming Binary Ionic compounds name of metal (cation) + name of anion NaCl sodium chloride CaO calcium oxide Cu2O copper(I) oxide cuprous oxide CuO copper(II) oxide cupric oxide Naming Polyatomic Ionic compounds BaCO3 barium carbonate Li2SO4 lithium sulfate Li2SO3 lithium sulfite

Covalent bonds Nonmetal-Nonmetal Metalloid-Nonmetal Sharing of valence electrons

Lewis Dot Structure H He Li Al C N Cl H Or Cl Lewis Structure Cl: 1s2 2s2 2p6 3s2 3p5 H: 1s1 He: 1s2 Ar: 1s2 2s2 2p6 3s2 3p6

Unshared pair of electrons (Lone pair) Cl H x Shared pair of electrons (bonding pair of electrons) Only valance electrons are involved in bonding (ionic and covalent bonds).

Electronegativity Electronegativity Ionization energy A measure of an atom’s attraction for the electrons Electronegativity Ionization energy

Nonpolar covalent bond: electrons are shared equally. Covalent bonds Nonpolar covalent bond: electrons are shared equally. Polar covalent bond: electrons are shared unequally. H Cl δ+ δ- Dipole

Electronegativity Difference Between Bonded Atoms Electronegativity & bonds Electronegativity Difference Between Bonded Atoms Type of Bond Less than 0.5 Nonpolar Covalent 0.5 to 1.9 Polar Covalent Greater than 1.9 Ionic H H 2.1 – 2.1 = 0 Nonpolar covalent N H 3.0 – 2.1 = 0.9 polar covalent Na F 4.0 – 0.9 = 3.1 Ionic

Covalent compounds H H – CH4 H C H H – C – H – H H O O O O – CH2O – – H C H H – C – H H C H H – C – H Correct H H – NH3 H N H H – N – H

H H H H H C C H H C C H C2H4 H H H H C – C C = C H H H H Correct H C C H H C C H C2H2 H – C – C – H H – C C – H

Resonance When a molecule has more than one Lewis structure. _ CH3COO- Resonance structures Resonance structures _

Some exceptions to the Octet rule B H 6

Carbon: normally forms four covalent bonds and has no unshared pairs of electrons. Hydrogen: forms one covalent bond and no unshared pairs of electrons. Nitrogen: normally forms three covalent bonds and has one unshared pair of electrons. Oxygen: normally forms two covalent bonds and has two unshared pairs of electrons. A Halogen: normally forms one covalent bond and has three unshared pairs of electrons. C H N . . . . . O = . . . . . Cl . .

electronegative element + ide Naming Binary Covalent compounds Mono – Di – Tri – Tetra – Penta – Hexa – Hepta – Octa – Nona – Deca Don’t use “mono” for the 1st element. Drop the “a” when followed by a vowel. prefixes and name the less electronegative element + prefixed and the name the more electronegative element + ide NO2 nitrogen dioxide N2O4 dinitrogen tetroxide CCl4 carbon tetrachloride S2O3 disulfur trioxide

VSEPR: Valence-Shell Electron-Pair Repulsion method VSEPR Model VSEPR: Valence-Shell Electron-Pair Repulsion method Bond angle: angle between two atoms bonded to a central atom. Regions of electron like to be as far away as possible from the others.

Regions of electron density Four regions of electron density around an atom:

Bond Angles & Geometric Structures Linear molecules 2 regions Trigonal planar molecules 3 regions Tetrahedral arrangement 4 regions

Tetrahedral Electron Pair Geometry (Tetrahedral arrangement) Methane (CH4) Ammonia (NH3) Water (H2O) Tetrahedral Trigonal Pyramidal Bent (V Shaped)

Tetrahedral Electron Pair Geometry Unshared electron paires H2O CH4 NH3

Centers of δ+ and δ- lie at different places (sides). Polarity It has polar bonds. Centers of δ+ and δ- lie at different places (sides). H δ+ – δ- δ- δ+ δ+ δ+ H – C – H O = C = O – δ- H δ+ nonpolar molecule δ- O δ- = N C H H δ+ H H δ+ H polar molecule