1. Review of Chapter 6: Bonding

2.  Bonds are forces of attraction between (-) electrons of one atom and the (+) nucleus of another atom, with 2 electrons in every bond  Forming bonds releases energy (exothermic)  Breaking bonds requires energy to be absorbed (endothermic)

3.  Valence electrons are the outermost electrons, farthest from the nucleus, involved in bonding  Elements in same column/group have same valence #  Right-most number listed in electron configuration on periodic table tiles  Octet Rule: Most atoms want to have 8 valence electrons, and make bonds to gain a share of 8  H and He only want 2 total  Noble gases already have 8, don’t react with others

4. Lewis Structures  Show # of valence electrons (1 to 8) around symbol for atom  Valence electrons drawn as dots

5.  Electronegativity is how strongly an atom wants electrons and pulls on them in a bond  Listed on Table S  Nonpolar covalent bonds: difference between EN values of two bonded atoms is 0  Polar covalent bonds: difference between EN values of two bonded atoms is between 0.1 and 1.7  Ionic bonds: difference between EN values of two bonded atoms is greater than 1.7

6. Covalent Bonding  Nonmetals share electrons to form complete octets  Any molecule with ONLY NONMETALS is covalently bonded  Can be single/double/triple bonds, with extra electrons as lone pairs surrounding atoms

7. 1. Count total # valence electrons in atoms of compound 2. Arrange atoms Central usually has lowest electronegativity, only once, isn’t H 3. Place single bonds 4. Add lone pairs to outsides, then center 5. Make more bonds if needed 1. Count total # valence electrons in atoms of compound 2. Arrange atoms Central usually has lowest electronegativity, only once, isn’t H 3. Place single bonds 4. Add lone pairs to outsides, then center 5. Make more bonds if needed

8.  Ions form when atoms lose or gain electrons  Ions are just atoms with (+) or (-) charges  Metals lose e-, becoming (+) cations  Ionization energy is energy needed to take an electron away from an atom to make an ion  Nonmetals gain e-, becoming (-) anions  In ionic bonding, metals give electrons away to nonmetals; charged ions then hang out near each other, attracted by different charges Ionic Bonding

9.  To draw Lewis structures for metal and nonmetal ions: When ions form, metal gets no e-, nonmetal gets a complete set of 8. Put each ion in brackets, and write the charge at top right (oxidation number from periodic table)

10. Metallic Bonding  Inside pieces of metal, (+) charged metal atoms are lined up neatly, with (-) electrons in constant motion, moving throughout the whole structure  This “sea of mobile electrons” holds the metal together, makes it a good electrical conductor, and makes it malleable (easily shaped)

11. Intermolecular Forces  Non-permanent, not “real bonds”  Influence melting and boiling points  Electrostatic attractions between different charges  Hydrogen bonding: temporary attraction between an H atom and an atom of either F/N/O  Responsible for many important properties of water