Covalent Bonding and Naming

1. Covalent Bonding and Naming Chemistry 11 Mrs. Kay Read Pages 168-171, 185-196

2. 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 (H2)

3. HOBrFINCl • These elements naturally form as diatomic molecules (2 atoms bonded covalently) • Hydrogen (H2), Oxygen (O2), Bromine (Br2), Fluorine (F2), Iodine (I2), Nitrogen (N2), and Chlorine (Cl2)

4. 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, F2

5. 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

6. Triple Bonds • Sharing 6 electrons, 3 pairs of electrons • Shortest and strongest of the covalent bonds • Ex: nitrogen, N2

7. 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: P2O4 = diphosphorous tetroxide

8. Practice: • CO2 • NH3 • BF3 • NO2 • N2H4 • N2F2 • Carbon dioxide • Nitrogen trihydride • Boron trifluoride • Nitrogen dioxide • Dinitrogen tetrahydride • Dinitrogen difluoride

9. 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

11. 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)

12. 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 = 3.04-2.20 = 0.84 (more polar) H-C: ΔEN = 2.55 – 2.20 = 0.35

13. 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

14. 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)

15. Lewis Structures • Look at valence electrons of all atoms • Pick a central atom (least electronegative usually, has most bonding sites) • 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)

16. 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)

17. Some examples

18. Practice drawing and naming Lewis Structures • H2O • CH2O

20. 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 [CO3]2- and [NH4+]

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