Chemical Bonding Notes

Valence electrons are used in bonding. Stable elements want to achieve 8 electrons similar to the noble gases A metal wants to achieve the configuration for the noble gas before. A nonmetal wants to achieve the configuration for the noble gas after.

Bond- force that holds groups of two or more atoms together and makes them function as a unit Bond energy- energy required to break the bond (tells the bond strength) Ionic bonding- between ionic compounds which contain a metal and a nonmetal Atoms that lose electrons relatively easily react with an atom that has a high affinity for electrons Transfer of electrons Covalent bonding- between two nonmetals Electrons are shared by nuclei Polar Covalent bonding- unequal sharing of electrons positive end attracted to the negative end  (delta) indicates partial charge

Electronegativity-(p. 362) relative ability of an atom in a molecule to attract shared electrons to itself The higher the atom’s electronegativity value, the closer the shared electrons tend to be to that atom when it forms a covalent bond Increases – across a period Decreases- down a group Electronegativity difference Bond type Zero ( )Covalent Intermediate (0.4 – 1.4) Polar covalent Large (>1.4)Ionic

Bond Type Ex. Determine if each bond is covalent, polar covalent or ionic. H-H, O-H, Cl-H, S-H, F-H H-H = = 0 covalent O-H = = 1.4 polar covalent Cl-H = =.9 polar covalent S-H = =.4 covalent F-H = = 1.9 ionic

Naming Molecular Compounds Prefixes: (MEMORIZE) Mono-1tetra-4hepta-7deca-10 di-2penta-5octa-8 tri-3hexa-6non-9 prefixes are used with both the first named and second named element. Exception: mono- is not used on the first word second word ends in –ide If a two syllable prefix ends in a vowel, the vowel is dropped before the prefix is attached to a word beginning with a vowel monooxide Writing Molecular Formulas Translate prefixes Molecule (molecular compound) – term used to describe covalent compounds, made from two nonmetals

N 2 Odihydrogen monoxide Si 8 O 5 tetrasulfur hexachloride NH 3 carbon monoxide P 3 I 10 carbon dioxide = Dinitrogen monoxide = Octasilicon pentoxide = Nitrogen trihydride = Triphosphorus deciodide = H 2 O = S 4 Cl 6 = CO = CO 2 Molecule Names & Formulas Examples:

Lewis Structure- representation of a molecule Shows how the valence electrons are arranged among the atoms in the molecule. s p z X p x p y Oxygen 1s 2 2s 2 2p 4 For an element: O For a compound: Li +  [Li] +1 + [ Cl ] -1 Cl For a molecule: F

Duet rule- He & He only need 2 electrons for a full, stable valence shell Octet rule- eight electrons needed for a full, stable valence shell Bonding pair- electrons shared with other atom Lone pair - not involved in bonding Line (─) = 2 shared electrons dots () = 2 unshared electrons

5 Steps for Covalently Bonded Lewis Structures 1.Count to total number of valence electrons. 2.Arrange the atoms (least electronegative in center – except H & O) 3.Give each atom a single bond to the center atom 4.Distribute the remaining electrons from the total as lone PAIRS 5.Check that each atom meets the octet rule –Exceptions: Be – only 4e -, B – only 6e -, P – up to 10e-, S – up to 12e - –Not meeting to octet rule? move a lone pair to create a double bond. –Still not meeting the octet rule? move a second lone pair to create a triple bond. Ex. GeBr 4 Valence = 1(4) + 4(7) = 32 Ge Br

Single bond- involves two atoms sharing one pair Double bond- involves two atoms sharing two pairs Triple bond- involves two atoms sharing three pairs Ex. CH 4 C 2 H 4 C 2 H 2 Valence: 1(4) + 4(1) = 8 Valence: 2(4) + 4(1) = 12 Valence: 2(4) + 2(1) = 10 C H H H H CC H H H H CCHH

Resonance- more than one Lewis structure can be drawn for the molecule Ex. CO 2 Valence:1(4) + 2(6) = 16 COO COO COO