Chapter 11: Chemical Bonding Chemistry 1020: Interpretive chemistry Andy Aspaas, Instructor

Chemical bonds Chemical bonds: forces that hold atoms together Ionic bonds: forces of attraction between oppositely charged ions –Electron transfer: forms two oppositely charged ions –Electrostatic forces: opposite charges attract Covalent bonds: forces of attraction between two atoms which are sharing electrons –Molecule: group of covalently bonded atoms

Charges on ions Ionic compound: combination of oppositely charged metal and nonmetal ions Stable ions form noble gas-like electron configurations –Metals form cations: Cations have lost electrons to get noble gas configuration –Nonmetals form anions: Anions have added electrons Use charges to predict ionic formulas –Ionic compounds have no net charge

Covalent bonds Covalent bond: shared pair of electrons –Hold nonmetal atoms together in a molecule Polar covalent bond: covalent bond between unlike atoms –Unequal sharing of electrons –One end of bond has larger electron density than other Bond polarity: result of uneven electron sharing –End with larger electron density gets partial negative charge (  -) –End that is electron deficient gets partial positive charge (  +)

Electronegativity Electronegativity: ability of an atom to attract shared electrons –Values from –Large values: atom attracts electrons more strongly Periodic table: electronegativity increases left to right (across a period) –Decreases top to bottom (down a group) Larger differences in electronegativity between covalently bonded atoms mean a more polar bond

Dipole moment If a molecule has a center of positive charge and a center of negative charge in different points, it has a dipole moment If there are more than one partial negative or positive charges in a molecule, they may partially cancel each other out –Combine to form a single dipole moment for the molecule Molecules with a large dipole moment are polar

Lewis structures Chemical bonding involves only valence electrons of atoms Lewis structure: shows valence electrons as dots around atoms –Cations have no dots, anions have 8 Duet rule: Hydrogen forms stable molecules when it shars two electrons Octet rule: Second-row nonmetals form stable molecules when valence orbitals are full, 8 electrons

Writing Lewis structures Find sum of all valence electrons in molecule Use one pair of electrons to connect each pair of bound atoms Arrange remaining elecctrons to satisfy the duet rule for hydrogen or the octet rule for 2nd-row elements

Multiple-bonds It’s possible for a pair of atoms to share 2 or 3 pairs of electrons in order to satisfy the octet rule –Double bond: 2 pairs of electrons are shared –Triple bond: 3 pairs of electrons are shared

A few exceptions to the octet rule Boron compounds are stable with 6 electrons in boron’s valence shell –BF 3 Some third row elements can expand their octet –S, P

3-dimensional molecular structure VSEPR: valence shell electron pair repulsion model –3-dimensional molecular structure is determined by minimizing repulsions between electron pairs –Count electron pairs as well as bonds

Electron pair arrangements Linear: only 2 total lone pairs and/or bonds –180° bond angles –BeCl 2 Trigonal planar: 3 total lone pairs and/or bonds –120° bond angles –BF 3 Tetrahedral: 4 total lone pairs and/or bonds –109.5° bond angles –CH 4

Molecule shape Arrangement of bonds indicate shape of molecule 3 bonds + 1 lone pair –Trigonal pyramid –NH 3 2 bonds + 2 lone pairs –Bent –H 2 O