Chemical Bonding

A chemical bond is a mutual attraction between nuclei and valence electrons of different atoms that binds atoms together. As single particles, atoms have a high potential NRG. Nature favors a reduction in this energy-by bonding, they reduce their potential NRG. Why do atoms form bonds?

The Octet Rule Remember….for atoms to be most stable (i.e. have the lowest possible potential NRG) they must have their outermost shell full of electrons. The Octet Rule states: Atoms require 8 valence electrons to be stable ….There are exceptions to the octet rule. For example: H and He only need 2 electrons. Boron needs six and Beryllium only needs four.

Electron-Dot Notation Used to show the number of valence electrons an atom

Periodic Table w/ Electron Dot Notations

Types of Bonds Covalent Ionic Metallic

Covalent Bonding Results from the sharing of electron pairs between two atoms. Groups of atoms held together by covalent bonds are called molecules.

Properties of Covalent Bonds Low melting and boiling points Form between two nonmetals Form by the sharing of electrons Often liquids or gases at room temp. Will either be nonconductors or poor conductors of electricity

Bonds: Single bonds- one pair of electrons shared Ex. Cl and Cl Double bonds- sharing of two pairs of electrons Ex. Ethene (C2H4) Triple bonds- sharing of three pairs of electrons Ex. N and N

Lewis Structures Shows electrons being shared in covalent bonds as well as any unshared electrons (also called a lone pair) that are not involved in bonding. Examples: H and H N and N C and O

Covalent Bonds can be either: Polar Results from unequal sharing of electrons Nonpolar Results from equal sharing of electrons

Differences in electronegativity determine if a covalent bond will be polar or nonpolar-the bigger the difference-the more likely it will be polar

Predicting Bond Type Electronegativity Diff. Bond Type 0-0.3 Nonpolar Covalent 0.3-1.7 Polar Covalent 1.7-3.3 Ionic Bond

Predict the Bond Type: Bonding between sulfur and Hydrogen Cesium Chlorine

Intermolecular Forces Three types of intermolecular forces: Dipole-Dipole (only present in polar molecules) Hydrogen bonding (only present in polar molecules) London dispersion forces (present in both polar and nonpolar) These are forces that occur between COVALENTY BONDED ELEMENTS

Dipole-Dipole The strongest intermolecular forces occur between polar molecules because they act as tiny dipoles (created by equal but opposite charges separated by a short distance) These oppositely charged areas are attracted to one another.

Hydrogen Bonding Intermolecular force in which a H atom is bonded to a highly electronegative atom This is a particularly strong type of dipole-dipole force

London Dispersion Forces Any atom or molecule will experience weak intermolecular forces due to the continuous motion of electrons. This movement creates uneven electron distribution which creates a temporary dipole. London dispersion force- an attraction resulting from uneven electron distribution and the creation of temporary dipoles These are very weak forces.

Remember….. In covalent bonds: Polar bonds are stronger than nonpolar In reference to intermolecular forces: H-bonding is strongest, then dipole-dipole, then dispersion forces

Boiling Points and Intermolecular Forces Boiling points can give information about bond strength. Stronger bonds have higher boiling points and weaker bonds have lower boiling points. Examples: H2 is nonpolar covalent and boils at -253 C NH3 is polar covalent and boils at -88 C NaCl is ionic and boils at 1413 C

Molecular Geometry VSEPR theory (valence-shell, electron pair repulsion) Repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible

Molecules w/no unshared pairs that form linear molecules. Formula Lewis Dot BeF2 H2 HCl Will have a sample formula of AB2.

Molecules w/ no unshared pairs that form trigonal planar molecules. Formula Lewis Dot BCl3 Will have a sample formula of AB3.

Molecules w/no unshared pairs that form tetrahedral molecules. Formula Lewis Dot CH4 Will have a sample formula of AB4.

Molecules w/unshared pairs of electrons. VSEPR states that unshared pairs (lone pairs) of electrons will occupy a space around the atom just as bonding pairs do. Example: water (H2O) The unshared pair of electrons on the oxygen atom results in a “bent” molecule. Has the formula AB2E2

More examples: Formula Lewis Dot NH3 Has the formula AB3E

PHET Simulations

sp3 Hybridization Used to explain how the orbitals of an atom become rearranged when the atom forms covalent bonds Hybridization is the mixture of two or more atomic orbitals of similar energies on the same atom to produce hybrid orbits of equal energy Let’s look check out methane, CH4

Geometry of Hybrid Orbitals Hybridization # of Hybrid Orbitals Geometry s, p sp 2 Linear s,p,p sp2 3 Trigonal planar s,p,p,p sp3 4 tetrahedral

Ionic Bonding Results from the electrical attraction between a cation and an anion Bonds are formed when electrons are transferred. By transferring electrons-their outermost shells become filled.

Two Types of Ions Monatomic -ions that form from one atom Ex. Na+1 Cl-1 Mg+2 S-2 Polyatomic Ions that form from more than one atom Ex. SO4-2 (sulfate) NH4+1 (ammonium)

Oxidation Numbers Groups 1-18 Group Oxidation # 1 +1 2 +2 13 +3 14 +4 or -4 15 -3 16 -2 17 -1 18 0 As electrons are gained or lost, the atom takes on either a negative or positive charge which is called an oxidation number.

Characteristics of Ionic Bonds High melting pts. High boiling pts. Hard and brittle can conduct electricity in solution Form between a metal and a nonmetal

Metallic Bonding Chemical bond that results from attraction between metal atoms and the surronding “sea” of electrons created by the overlapping of orbitals where electrons can move freely Characteristics: Good conductors of heat and electricity Absorb a wide range of light frequencies Malleable Ductile Form between two metals