Covalent Bonding and Naming Chemistry 11 Mrs. Kay Read Pages ,

Pure Covalent Bonding  Equal sharing of electron between two of the same non-metals  The electronegativity between the two atoms must ZERO!  Ex: hydrogen gas (H 2 )

HOBrFINCl  These elements naturally form as diatomic molecules (2 atoms bonded covalently)  Hydrogen (H 2 ), Oxygen (O 2 ), Bromine (Br 2 ), Fluorine (F 2 ), Iodine (I 2 ), Nitrogen (N 2 ), and Chlorine (Cl 2 )

Single Bonds  Sharing of 2 electrons, one pair  A single line represents the 2 electrons  Longest and the weakest of the covalent bonds (easiest to break apart)  Ex: Fluorine, F 2

Double Bonds  Sharing 4 electrons 2 pairs of electrons  Shorter and stronger than a single bond, takes more energy to break it apart.  Drawn with two lines, each line represents 2 electrons

Triple Bonds  Sharing 6 electrons, 3 pairs of electrons  Shortest and strongest of the covalent bonds  Ex: nitrogen, N 2

Naming simple molecules  If its diatomic (HOBrFINCl) you simply name the non- metal its made of Must memorize the prefixes  RULES: if there is only one of the first atom than don’t use a prefix, otherwise use a prefix.  Ex: CO = carbon monoxide  Ex: P 2 O 4 = diphosphorous tetroxide PrefixNumber Mono1 Di2 Tri3 Tetra4 Penta5 Hexa6 Hepta7 Octa8 Nona9 Deca10

Practice: 1. CO 2 2. NH 3 3. BF 3 4. NO 2 5. N 2 H 4 6. N 2 F 2 1.Carbon dioxide 2.Nitrogen trihydride 3.Boron trifluoride 4.Nitrogen dioxide 5.Dinitrogen tetrahydride 6.Dinitrogen difluoride

Polar Covalent Bonding  Electrons are shared unequally.  They are not ionic, because the electron is not totally removed because there was not enough attraction to totally remove the electron.  Based on difference in electronegativity  Ex: HCl

Electronegativity  The degree to which an atom attracts electrons to itself  It is not a measurement, but a scale.  Periodic trend: generally increases from left to right across a period and from bottom to top in a group.  What is the most electronegative element?  Fluorine (F)

 The greater the electronegative difference, the more polar the bond because the more electronegative atom will attract the “shared” electron pair closer to itself. Ex: H-N = more polar covalent than H-C, because electronegativity of H (2.20) and N (3.04), while C (2.55) H-N: ΔEN = = 0.84 (more polar) H-C: ΔEN = 2.55 – 2.20 = 0.35

Bonding Continuum  Idea that bonds can behave as mostly ionic or mostly covalent  Use the difference in electronegativity to label the bond type  0 = pure covalent  0.4 to 1.7 = mostly covalent  1.7 or greater = mostly ionic

Lewis Structures for molecules  Need to show the structure of a molecule.  Will use Lewis structures (electron dot diagrams) to show where there are lone pairs (filled orbitals) and bonding pairs (places where bonds most likely occur)

Lewis Structures 1. Look at valence electrons of all atoms 2. Pick a central atom (least electronegative usually, has most bonding sites) 3. Align all atoms so that each have their ideal amount of valence electrons achieved through sharing. Usually 8 (stable octet), but can be 2 (H, He) and 6 (B)

Carbon tetrachloride  Carbon is the central atom.  It has 4 bonding pairs.  Chlorine wants to share one bonding site each.  Need 4 chlorines for every one carbon (Cl has 3 lone pairs and 1 bonding pair)

Some examples

Practice drawing and naming Lewis Structures  H 2 O  CH 2 O

What about ions?  Count up all valence electrons that you are allowed to place.  Still pick the central atom.  Still have the correct number of electrons around each atom (usually 8, except for H and He)  Add extra electrons if an anion and take away electrons if a cation TRY [CO 3 ] 2- and [NH 4 + ]

Metallic Bonding What is a Metallic Bond? - A metallic bond occurs in metals. A metal consists of positive ions surrounded by a “sea” of mobile electrons. Name 4 Characteristics of a Metallic Bond. 1.Good conductors of heat and electricity 2.Great strength 3.Malleable and Ductile 4.Luster This shows what a metallic bond might look like.