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Chemistry Study Guide #4

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

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