Chapter 11: Chemical Bonding

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

2. 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

3. 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

4. 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 (+)

5. Electronegativity • Electronegativity: ability of an atom to attract shared electrons • Values from 0.7 - 4.0 • 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

6. 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

7. 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

8. 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

9. 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

10. A few exceptions to the octet rule • Boron compounds are stable with 6 electrons in boron’s valence shell • BF3 • Some third row elements can expand their octet • S, P

11. 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

12. Electron pair arrangements • Linear: only 2 total lone pairs and/or bonds • 180° bond angles • BeCl2 • Trigonal planar: 3 total lone pairs and/or bonds • 120° bond angles • BF3 • Tetrahedral: 4 total lone pairs and/or bonds • 109.5° bond angles • CH4

13. Molecule shape • Arrangement of bonds indicate shape of molecule • 3 bonds + 1 lone pair • Trigonal pyramid • NH3 • 2 bonds + 2 lone pairs • Bent • H2O