
11/5/03
Ions- Atoms that have gained or lost one or more electrons. When p+ doesn't equal e-.
Cations- Positively charged ions. When p+ > e-.
Anions- Negatively charged ions. When e- > p+.
Writing Ionic Formulas
- Cation first, anion second
- Use charges (yellow sheet) to write lowest whole number ratio possible
o Ex. Na (+1) and P (-3) is Na(3)P
- When naming, give name of cation, then the name of the anion with "ide" on the end. When the anion is polyatomic (parenthesis) don't add "ide". Don't forget roman numerals when the cation exists in multiple oxidation states.
11/10/03
Electoneutrality- Uncharged. When p+ = e-.
Isoelectronic- ion has same electron configuration as a different type of atom (due to gain/loss of electrons and trying to become like a noble gas)
Octet rule- if s and p orbitals are filled, the atom is stable
Ionic formulas- cation and anion bond in whole- number ratios
Ionic Bond- coulombic force between two objects that have electric charges
Halide- a salt made of cations and an anion of the halogen group
Electron transfer is usually endothermic
Characteristics of Ionic Compounds
- Solids at room temperature
- High melting and boiling points (generally)
- Don't conduct electricity in solid form
- Hard and brittle
- Good electroconductors in liquid state and when dissolved in water (this breaks the compound apart)
Electolites- ions in liquid
Book
Oxidation #- describes distribution of electrons in bonded atoms
Minerals- inorganic crystalline materials found in rocks
Silicates- minerals with anions of lots of connected silicon and oxygen atoms
Hydrates- Ionic compounds with water molecules incorporated in their crystal lattices
Anhydrous- result of heating hydrate to remove water
11/12/02
valence electrons- outermost electrons
Ionic Bonds- use charges (yellow sheet) give and take e-.
Covalent Bonds- share e-. Two or more valence e- attracted to 2 atom's nuclei
Ionic compounds are formed between metals and nonmetals
Covalent compounds are formed between two nonmetals
Molecular orbit- A region where an electron pair is most likely to exist as it travels in (3D) space
Diatomic- weak covalent bond between two atoms of the same type. Found in H, N, O, Fl, Cl, Br, I.
Molecule- two or more atoms held by covalent bonds
Compound- two or more atoms held by ionic bonds
Nonpolar Covalent
- Equal sharing
- Two identical nonmetals
- EN (0 to .5)
Polar Covalent
- Unequal sharing (pull)
- Two different nonmetals
- EN (.5 to 2.1)
Ionic Bonding
- Give and take of electrons
- Metal and nonmetal
- EN (2.1 to 3.3)
11/14/03
Covalent Bonding Rules
- Highest electronegative element is second
- First element, name with full name
o Use prefixes, determined by subscripts (little yellow sheet)
o Never use Mono- on the first compound
- Last element, name using the prefix determined by subscripts and add "ide".
Lewis Structure 11/17/03
Gilbert Newton Lewis- 1875-1946- Discovered heavy water and came up with Lewish Structures
Octet Rule- every atom wants 8 electrons (fill s (2) and p (6) orbitals)
- exception is Hydrogen, which just wants two (to become like Helium)
Valence Electrons- bonding electrons of an atom
- Refer to a group number
Group 1 2 13 14 15 16 17 Valence Electrons 1 2 3 4 5 6 7
Unshared Pair- a pair of valence electrons not involved in bonding
Single Bond- sharing of two electons forming only one bond
Multiple bonds- sharing four or more electrons
Book 6-4
Intermolecular Forces- forces that exist between molecules
Dipole force- the force that exists with positive and negative ends of a polar molecule
Hydrogen bond- very strong dipole force
London forces- aka dispersion forces- overshadowed by dipole forces. Boiling point increases as number of electrons in molecules or atom increases