1. Chapter 4
Covalent Compounds
Dalton Trans., 2016,45,

2. Covalent Compounds
Covalent Compounds: compounds composed of atoms bonded to each other through the sharing of electrons
Electrons NOT transferred
No + or – charges on atoms
Non-metal + Non-metal
Also called “molecules”
Examples:
CO2
Cl2
CH4

3. Dihydrogen Monoxide (DHMO)!!!!!!
Properties – Deadly, colorless, odorless liquid
Hazards!!!
Can mutate DNA!!
Death due to accidental inhalation of DHMO, even in small quantities.
Prolonged exposure to solid DHMO causes severe tissue damage
DHMO is a major component of acid rain
Contributes to soil erosion.

4. Dihydrogen Monoxide H2O, commonly known as water
Can mutate DNA!!— water can get ionized to react with DNA
Death due to accidental inhalation of DHMO, even in small quantities--- when someone drowns
Prolonged exposure to solid DHMO causes severe tissue damage--- Ice
DHMO is a major component of acid rain
Contributes to soil erosion--- water corrodes the landscape over time

5. Covalent Bonding
or H-H
Duet or

6. Naming Covalent Compounds
Name the first non-metal by its elemental name
Add a prefix to indicate how many
If only one atom, don’t put mono 1
mono 6
hexa 2 di 7
hepta 3
tri 8
octa 4
tetra 9
nona 5
penta 10
deca
4) Name the 2nd non-metal and change its ending to “-ide”
5) Add a prefix to indicate how many

7. Problems Write the name of the following compounds: CO NI3 N2O SF6
B2O3
carbon monoxide
nitrogen triiodide
dinitrogen monoxide
sulfur hexaflouride
diboron trioxide

8. Write the formula for the following compounds:
Phosphorous Pentachloride
Nitrogen Monoxide
Dinitrogen Tetroxide
Tetraphosphorous Decoxide
PCl5 NO
N2O4
P4O10

9. Problems KCl Na2S H2O SO2 K3PO4 FeCl3 (NH4)2SO4 SCl2 Cu(OH)2 P2O5
Potassium chloride
Sodium sulfide
Dihydrogen monoxide
Sulfur dioxide
Potassium phosphate
Iron(III) chloride
Ammonium sulfate
Sulfur dichloride
Copper(II) hydroxide
Copper(II) hydroxide

10. Phosphorous Pentabromide Magnesium Nitride
Sodium Iodide
Aluminum Sulfate
Phosphorous Pentabromide
Magnesium Nitride
NaI
Al2(SO4)
PBr5
Mg3N2

11. Naming Acids Acids that do not contain oxygen
Begin the name with “hydro”
Name the anion, but change the ending to “-ic”
Add “acid” on the end
HCl = hydrochloric acid
HF = hydrofluoric acid
EXCEPTION, if in the gas phase, treat like a regular covalent compound for naming with no prefixes
HCl(g) = hydrogen chloride
Naming Acids

12. Acids that contain oxygen/oxyanions
Do not put “hydro” at the beginning
Begin the name with the anion
If the anion has the ending “-ate,” change this to “-ic acid”
If the anion has the ending “-ite,” change this to “-ous acid”
HClO4 perchloric acid
HClO3 chloric acid
HClO2 chlorous acid
HClO hypochlorous acid

13. Problems Name the following HBr(g) HBr(aq) Hydrogen bromide HNO2(aq)
HI (aq)
HI (g)
H2CO3 (aq)
H3PO4 (aq)
H3PO3 (aq)
HCN (aq)
Hydrogen bromide
Hydrobromic acid
Nitrous acid
Nitric acid
Hydroiodic acid
Hydrogen iodide
Carbonic acid
Phosphoric acid
Phosphorous acid
hydrocyanic acid

14. Molecular Structures

15. Ball & Stick Models
Space-Filling Models
Water H2O
Methane
CH4

16. Octet Rule
or H-H
Duet or

17. Making Lewis Dot Structures
Count the total number of valence electrons in the molecule.
Ex: PCl3
2) Use atomic symbols to draw a proposed structure with shared pairs of electrons
Pick a central atom – atom that wants to make the most bonds
Atoms don’t tend to bond to other atoms of the same element when they can avoid it
Exception: Carbon

18. Place lone pair electrons around each outside atom (except H) to satisfy the octet rule, beginning with the terminal atoms
Place any leftover electrons on the central atom
If the number of electrons around the central atom is less than 8, change single bonds to the central atom to multiple bonds (double or triple).
Ex: CH2O

19. Covalence chart Atom Covalence number Non bonding electrons H 1
C or Si 4
P or N 3
O or S 2
F, Cl, Br, or I B
Covalence chart

20. What Certain Atoms Like To Do
Halogens
Like to have one single bond and 3 lone pairs (non-bonding electrons)
F,Cl, Br, I
Carbon
Likes to have 4 single bonds and no lone pairs
A double bond counts as two singles
A triple bond counts as three singles
Likes to be central
Likes to bond to other carbons

21. Silicon
Likes to do what carbon does
Oxygen
Likes to have two single bonds and 2 lone pairs
Sulfur
Likes to do what oxygen does
May expand its octet
Nitrogen
Likes to have 3 single bonds and one lone pair

22. Phosphorous
Likes to do what nitrogen does
May expand its octet
Hydrogen
Likes to be terminal with only one single bond
No lone pairs!
Boron
Likes 3 bonds and no lone pairs (sextet)

23. Draw the LDS’s for the following molecules:
Cl2O
C2H4
C2H6O
Problems

24. Problems Draw the Lewis Structures for the following molecules: SH2
C3H8
Si2H6
PI3
CH3OH
C2H2
BF3
CCl2O

25. N2H4
CH2OS
C2H6O P4
C6H6

26. Dot structures for ions
Polyatomic ions contain covalent bonds
Covalent chart can be violated, but the octet rules is not
Comparing current e-’s to expected valence e-’s can tell you “charge” of an atom
One less e- = positive charge
One more e- = negative charge

27. Resonance
Two or more electron arrangements are possible for the same arrangement of atoms
The structures differ in only the multiple bond and lone pair placement

28. Resonance meaning
The true structure is an average of all resonance structures
Resonance allows electros to move (spread) across multiple atoms, this stabilizes the molecule as a whole
Spread out electrons aka “delocalized” electrons

29. Practice with resonance
N2O

30. The measure of the ability of an atom to attract electrons to itself
Increases across period (left to right) and
Decreases down group (top to bottom)
fluorine is the most electronegative element
francium is the least electronegative element
Electronegativity

32. Electronegativity Scale

33. Types of Bonding Non-Polar Covalent Bond:
Difference in electronegativity values of atoms is 0.0 – 0.4
Electrons in molecule are equally shared
Examples: Cl2, H2, CH4
ENCl = 3.0
= 0
Pure Covalent

34. Polar Covalent Bond:
Difference in electronegativity values of atoms is 0.4 – 2.0
Electrons in the molecule are not equally shared
The atom with the higher EN value pulls the electron cloud towards itself
Partial charges
Examples: HCl, ClF, NO
ENCl = 3.0
ENH = 2.1
3.0 – 2.1 = 0.9
Polar Covalent

35. Ionic Bond: Difference in EN above 2.0
Complete transfer of electron(s)
Whole charges
ENCl = 3.0
ENNa = 1.0
3.0 – 0.9 = 2.1
Ionic

37. Problems
Draw the lewis dot structure then predict the type of bonding in the following compounds using differences in EN values of the atoms. Indicate the direction of the dipole moment if applicable
KBr HF
BrI O2

38. Valence Shell Electron Pair Repulsion Theory
VSEPR theory:
Electrons repel each other
Electrons groups in a molecule arrange themselves so as to be as far apart as possible
Minimize repulsion
Determines molecular geometry

44. Defining Molecular Shape
Electron pair geometry: the geometrical arrangement of electron groups around a central atom
Atoms and lone pairs count as electron groups
Molecular Geometry: the geometrical arrangement of atoms around a central atom
Ignore lone pair electrons

45. 2 e- groups surrounding the central atom
e- pair geometry: linear
MG: linear
AXE designation: AX2E0
A: Central Atom
X: Bonding pairs
E: Non-bonding pairs
Example: BeCl2, CO2

46. 3 e- groups 3 Bonds, 0 Lone Pairs 2 Bonds, 1 Lone Pair
e- PG: Trigonal Planar (Triangular planar)
MG: Trigonal Planar
AX3E0
BF3
2 Bonds, 1 Lone Pair
e- PG: Trigonal Planar (Triangular planar)
MG: Bent/angular
AX2E1
GeCl2

47. 4 e- groups 4 bonds, 0 Lone Pairs 3 bonds, 1 Lone Pair
e- PG: Tetrahedral
MG: Tetrahedral
AX4E0
CH4
3 bonds, 1 Lone Pair
e- PG: Tetrahedral
MG: Triangular Pyramidal
AX3E1
NH3

48. 2 bonds, 2 Lone Pairs
e- PG: Tetrahedral
MG: Bent/Angular
AX2E2
H2O

51. Molecular Polarity

53. Drawing LDS With Correct Geometry
Commonly seen for tetrahedral geometry

54. Problems BF3 CH2O CBr4 CHCl3 CH2Cl2
Draw the 3D Lewis Dot Structures, using wedges and dashes when applicable, for the following molecules and then identify the net dipole, if any.
BF3
CH2O
CBr4
CHCl3
CH2Cl2

55. Chapter 4 checklist
Covalent compounds- what is a covalent compound? Know the rules for naming.
Lewis dot structures – how to draw molecules
Electronegativity – What is it and how does it determine the type of bond? Pure covalent, covalent, and ionic bonds depend on the electronegativity difference between the each atom in question
Valence shell electron pair repulsion theory
Resonance
Molecular Geometry and molecular polarity