Why do atoms form bonds? To attain a noble gas configuration How do atoms form bonds? By gaining, losing, or sharing electrons Gain or loss of electrons  ionic bonding Sharing of electrons  covalent bonding

Covalent Bonding Results from: electrostatic attraction between nucleus one atom & electrons of neighbor atom

The electrons are shared, not transferred Formed between 2 nonmetal atoms - sometimes two atoms of same element If attractions are > than repulsions: bond is formed

Potential Energy vs. Internuclear Distance Low energy High stability Bonding Systems

Different ways of representing a covalent bond Compounds with covalent bonds are molecular!

Recognizing covalent formulas Covalent bonding occurs between nonmetal atoms so… –formulas contain only nonmetals

Identify Bond Type from Formula ZnO N2O5N2O5N2O5N2O5Al CH 3 OH Al 2 O 3 CaBr 2 AuAg CO 2 Li 3 N Na 2 S Mg CsF H2OH2OH2OH2ONaCl SO 2 Cu CH 4 CovalentCovalent CovalentCovalent Covalent CovalentIonic Ionic Ionic IonicIonicIonic Ionic Metallic Metallic Metallic MetallicMetallic

Making a bond Liberates energy! Separate Atoms Molecule

A + B  AB A and B are both atoms AB is a molecule In this equation: –no bonds broken; one bond formed –Energy is released The reaction is exothermic

A + B  AB bond is formed energy is released AB has less PE than A + B AB is more stable than A + B

Breaking a Bond Absorbs Energy!

CD  C + D CD is a molecule; C and D are atoms In this equation: –one bond broken –no bonds formed This reaction is endothermic C & D have higher PE than CD

Structure of Covalent Compounds They form molecules The more reactants you have, the more molecules you can make

Representations of Molecules Space-filling model Ball-and-stick model Ball-and-stick and electron cloud

Representations of Molecules Lewis Structures of compound H 2 = Molecular Formula H-H = Structural Formula - represents 1 pair of shared electrons

Beyond H 2 : Other diatomic elements Hydrogen (H 2 ) Nitrogen (N 2 ) Oxygen (O 2 ) Fluorine (F 2 ) Chlorine (Cl 2 ) Bromine (Br 2 ) Iodine (I 2 ) These molecules are more stable than the individual atoms

Halogens Share 1 pair (2 electrons): form single covalent bond

Rules for drawing Lewis Diagrams 1.Arrange symbols on paper the way you think atoms are arranged - Diatomics are easy – they are next to each other - Hydrogens are always terminal - Atom with least attraction (lowest electronegativity) for shared electrons is placed in the center 2.Add up # valence electrons for all atoms starting with single bonds (electron pairs) between all atoms 3.Distribute electrons, starting with single bonds (electron pairs) between all atoms 4.Test validity (2 tests)

Tests for Lewis Structures Must Pass Both! 1.Number of dots = number of valence electrons found in step 2 2.Every atom has an octet of electrons around it (except H: only wants 2) Bonding electrons get counted 2 times – once for each atom sharing them

Assessing Lewis Diagrams If dot structure passes both tests, you’re finished If diagram fails one or both tests, try again If single bonds don’t work - try multiple bonds –Single bond = 2 electrons shared –Double bond = 4 electrons shared –Triple bond = 6 electrons shared

Try some examples H 2, F 2, HF, O 2, N 2 Step 1: Draw the symbols the way you think the atoms are arranged –Diatomics are easy – the atoms are right next to each other!

Drawing Lewis Diagrams H H Step 2: Count up the valence electrons valence electrons Each H has 1 valence electron so the total = 2 Step 3: Distribute the valence electrons, starting with single bonds between all atoms :

Testing Lewis Diagrams H H : Test 1: 2 dots in diagram = 2 valence electrons Test 2: Each H has 2 valence electrons (Remember H only wants 2)

Testing Lewis Diagrams F F Step 1 Step 2 2 X 7 = 14 valence electrons Step 3 F F : ::

Try HF H F Step 1 Step = 8 valence electrons Step 3 H F : :....

Types of Covalent Bonds Single: 2 atoms share 1 pair of electrons Double: 2 atoms share 2 pairs of electrons Triple: 2 atoms share 3 pairs of electrons

O O Step 1: Step 2: 2 X 6 = 12 valence electrons Step 3: O O : :: Step 4: Test failed No Good! Try O 2

Try O 2 Again! O O Step 1: Step 2: 2 X 6 =12 Step 3: Distribute electrons (single bond between atoms didn't work so try a double bond) :: Step 4: Test!

Try N 2 N N Step 1: Step 2: 2 X 5 = 10 electrons Step 3: Distribute electrons starting with a single bond between the nitrogen atoms :.. :.. : Step 4: Test No Good!

Try N 2 with a double bond N N :: : :.. No Good!

Try N 2 with a triple bond N N ::::: Step 4: Test