1. Classifying Chemical Compounds
Chemical Bonding
Classifying Chemical Compounds

2. Bond Types Atoms can combine to form compounds of two bond types:
Ionic
Compounds
Covalent
Compounds
Ex: NaCl
Ex: CO2

3. Physical Properties – Comparison
Property
Ionic
Compound
Covalent Compound
State at room temperature
Melting point
Electrical conductivity as a liquid
Solubility in water
Electrical conductivity when dissolved in water

4. Physical Properties – Comparison
Property
Ionic
Compound
Covalent Compound
State at room temperature
Crystalline solid
liquid, solid, gas
Melting point
Electrical conductivity as a liquid
Solubility in water
Electrical conductivity when dissolved in water

5. Physical Properties – Comparison
Property
Ionic
Compound
Covalent Compound
State at room temperature
Crystalline solid
liquid, solid, gas
Melting point
High
Low
Electrical conductivity as a liquid
Solubility in water
Electrical conductivity when dissolved in water

6. Physical Properties – Comparison
Property
Ionic
Compound
Covalent Compound
State at room temperature
Crystalline solid
liquid, solid, gas
Melting point
High
Low
Electrical conductivity as a liquid
Yes No
Solubility in water
Electrical conductivity when dissolved in water

7. Physical Properties – Comparison
Property
Ionic
Compound
Covalent Compound
State at room temperature
Crystalline solid
liquid, solid, gas
Melting point
High
Low
Electrical conductivity as a liquid
Yes No
Solubility in water
High solubility (for most)
Low solubility (for most)
Electrical conductivity when dissolved in water

8. Physical Properties – Comparison
Property
Ionic
Compound
Covalent Compound
State at room temperature
Crystalline solid
liquid, solid, gas
Melting point
High
Low
Electrical conductivity as a liquid
Yes No
Solubility in water
High solubility (for most)
Low solubility (for most)
Electrical conductivity when dissolved in water
Not usually

9. Chemical Properties – Comparison
Bonds form when the valence e- of atoms interact.
Atoms can either exchange e or share e-
There are two types of chemical bonds

10. Ionic Bonding When 2 atoms exchange (or transfer) electrons
1 atom loses an e-
1 atom gains an e-
metal non-metal
lose gain

11. Covalent Bonding When 2 atoms share electrons
Usually forms between 2 non-metals

12. Electronegativity Recall:
Electronegativity (EN) is a measure of an atom’s ability to attract electrons in a chemical bond.
When 2 atoms form a bond, each atom attracts the other’s electrons in addition to its own.

13. Table – Electronegativity Values

14. Predicting Bond Type Using EN
The difference between EN values (EN) for the 2 atoms in a bond can be used to predict bond type.
Ex: CaF2 EN = ?
Chemical bonds range from mostly ionic to mostly covalent.

15. Predicting Bond Type Using EN
EN values mostly polar mostly ionic covalent covalent

16. Predicting Bond Type Using EN - Examples
i) calcium fluoride EN =
CaF2
ii) hydrogen chloride EN =
HCl
iii) methane EN =
CH4
iv) oxygen EN = O2

17. Chemical Bonding
Bond Formation

18. Covalent Bonding Involves sharing of e-
Sharing of e- between atoms may not be equal.
non-polar covalent
(mostly covalent) bond involves equal sharing
polar covalent
bond involves unequal sharing

19. Covalent Bonding-Example
water (H2O) EN = 1.24
This is polar covalent bond (unequal sharing)
The O end of the H-O bond has a partial negative charge (-) while the H end of the bond has a partial positive charge ( +)
This separation of charges makes a dipole
( +   -)

20. Polar Bonds Using EN values you can predict the polarity of a bond.
The overall polarity of a molecule depends on a combination of:
the polarity of the bonds
the shape of the molecule

21. Shapes of Molecules
To illustrate the shape of the molecule using a Lewis diagram:
Determine the “central atom” , i.e. the atom requiring the most electrons.
Determine how many electrons each atom needs to gain in order to become isoelectronic with its closest noble gas.
Determine if multiple bonds are needed.
Determine the structural shape of the molecule based on
# atoms bonded to central atom
# lone pairs of electrons
(see handout of molecular shapes)

22. How Does That Look Again?
Shapes of Molecules

23. Shapes of Molecules-Examples
Chemical formula
Lewis diagram
Structural diagram
Shape
CH4
methane
CO2
H2O

24. Polarity of Molecules To predict the polarity of the molecule:
Determine if there are polar bonds using EN value
If there are polar bonds, establish the direction of the dipoles.
(A dipole exists when two opposite charges are separated by a short distance)
( +   -)

25. Polarity of Molecules iii. Once you identify the shape of the molecule
if the dipoles cancel due to the shape of the molecule then the compound is non-polar
if the dipoles do not cancel due to the shape of the molecule then the compound is polar

26. Polarity of Molecules-Examples
Chemical formula
Lewis diagram
Structural diagram
(include dipoles)
Shape
Polarity of molecule
(polar/
non-polar)
CH4
methane
CO2
H2O

27. Video Clip
Polar Bonding

28. Chemical Formulas - Ionic
Chemical Bonding
Chemical Formulas - Ionic

29. Writing Formulas–Binary Compounds
contain H, metal or metalloid with only 1 oxidation number (or valence)
(see periodic table for values)
name the first element and change the ending (suffix) of second element to “ide”
Use zero sum rule to write formula
Try Table 1

30. Chemical Formulas - Practice Table 1
Br1-
N3-
K1+
Mg2+
Al3+
K2S
KBr
K3N
potassium sulfide
potassium bromide
potassium nitride
Mg3N2
magnesium nitride
MgS
magnesium sulfide
MgBr2
magnesium bromide
Al2S3
aluminum sulfide
AlBr3
aluminum bromide
AlN
aluminum nitride

31. Tertiary Compounds contain polyatomic ions
(see periodic table given in class)
name the first element and then the polyatomic ion
Use zero sum rule to write formula
Try Table 2

32. Chemical Formulas - Practice Table 2
OH1-
CO32-
PO43-
Mg2+
Li1+
NH41+
hydroxide
carbonate
phosphate
Mg(OH)2
MgCO3
magnesium carbonate
Mg3(PO4) 2
magnesium phosphate
magnesium hydroxide
LiOH
lithium hydroxide
Li2CO3
lithium carbonate
Li3PO4
lithium phosphate
(NH4)2CO3
ammonium carbonate
NH4OH
ammonium hydroxide
(NH4)3PO4
ammonium phosphate
ammonium

33. Writing Formulas – Multivalent Ions
When the first element has more than 1 oxidation number (or valence)
IUPAC System (Stock System)
Use Roman numeral to indicate valence of more electronegative element
Ex:
Cu1+ copper (I)
Cu2+ copper (II)
Note: Hg2+ mercury (II)
Hg22+ (Hg1+)mercury (I)

34. Writing Formulas – Multivalent Ions
Old system(Classical System)(ic/ous)
Use suffix “ic” for highest valence and “ous” for lowest
Use classical name of element (see table provided on next slide)

35. Table: Multivalent Ion Names
Formula
IUPAC Name
Classical Name
Cu1+
copper (I)
Cu2+
copper (II)
Fe2+
iron (II)
Fe3+
iron (III)
Hg22+ (Hg1+)
mercury (I)
Hg2+
mercury (II)
Pb2+
lead (II)
Pb4+
lead (IV)
Sn2+
tin (II)
Sn4+
tin (IV)
Cr2+
chromium (II)
Cr3+
chromium (III)
Mn2+
manganese (II)
Mn3+
manganese (III)
Mn4+
manganese (IV)
Co2+
cobalt(II)
Co3+
cobalt(III)
Au1+
gold (I)
Au3+
gold (III)
Table: Multivalent Ion Names

36. Table: Multivalent Ion Names
Formula
IUPAC Name
Classical Name
Cu1+
copper (I)
cuprous
Cu2+
copper (II)
cupric
Fe2+
iron (II)
ferrous
Fe3+
iron (III)
ferric
Hg22+ (Hg1+)
mercury (I)
mercurous
Hg2+
mercury (II)
mercuric
Pb2+
lead (II)
plumbous
Pb4+
lead (IV)
plumbic
Sn2+
tin (II)
stannous
Sn4+
tin (IV)
stannic
Cr2+
chromium (II)
chromous
Cr3+
chromium (III)
chromic
Mn2+
manganese (II)
Mn3+
manganese (III)
Mn4+
manganese (IV)
Co2+
cobalt(II)
cobaltous
Co3+
cobalt(III)
cobaltic
Au1+
gold (I)
aurous
Au3+
gold (III)
auric
Table: Multivalent Ion Names

37. Chemical Formulas - Practice Table 3
IUPAC
Old System
FeCl2
FeCl3
copper (I) fluoride
nickel (II) sulfide
ferric oxide
mercuric oxide
iron (II) chloride
ferrous chloride
iron (III) chloride
ferric chloride
CuF
cuprous fluoride
NiS
nickelous sulfide
Fe2O3
iron (III) oxide
HgO
mercury (II) oxide

38. Writing Formulas – Polyatomic Ions
Some polyatomic ions contain the same types of atoms but different numbers of each type.
(see periodic table given in class)

39. Writing Formulas – Polyatomic Ions
For the related ions, only the “ate” ions need to be memorized, others may be found following this pattern:
Step method:

40. Writing Formulas – Polyatomic Ions
For the related ions, only the “ate” ions need to be memorized, others may be found following this pattern:
Step method:
As you move up one step…the # of oxygen atoms increases by 1
Charge on ion stays the same
hypo
ite
ate
per

41. Oxy-Acid Radicals Try the oxy-acid radicals handout provided
Note: Elements in the same group on the periodic table form the same polyatomic ion.

42. Chemical Formulas - Covalent
Chemical Bonding
Chemical Formulas - Covalent

43. Covalent Compounds-Prefix Method
When compounds contain 2 non-metals use the PREFIX METHOD for naming.
Write the name of the 1st element
the first element is the least electronegative element (i.e. closest to the LHS of the periodic table)
Write the name of the 2nd element deleting the last few letters and adding “ide”
Indicate the # atoms of each type with the appropriate prefix
See p105 Table 3.8 (copy this)
Note: “mono” is left out when there is a single atom of the first element.

44. Number
Prefix 1
mono- 2
di- 3
tri- 4
tetra- 5
penta- 6
hexa- 7
hepta- 8
octa- 9
nona- 10
deca-

45. Examples S2O3 N2O5 CF4 silicon monoxide diphosphorus tetroxide CO2 CO
N2O5
CF4
silicon monoxide
diphosphorus tetroxide
CO2
CO  
Al2O3
disulfur trioxide
dinitrogen pentoxide
carbon tetrafluoride
SiO
P2O4
carbon dioxide
carbon monoxide
aluminum oxide (IONIC - METAL!)

46. Hydrides These are binary compounds that contain Hydrogen
If H is less electronegative, then it is placed first in the name of the compound
hydrogen ion (cation) combines with non-metal but acts like a metal itself
H  H e-
e.g. H2Se hydrogen selenide
“metal” non-metal
HCl
hydrogen chloride

47. Hydrides
Sometimes H can act like an anion (usually with a Group 1 metal)… so it is placed second in the name of the compound
H + e-  H1-
e.g. NaH sodium hydride
CaH2
calcium hydride

48. Hydrides
Hydrogen containing compounds can also be formed with Family VA elements…but these are usually referred to by common names rather than IUPAC
e.g. NH3
PH3
AsH3
ammonia
phosphine
arsine

49. Acids
Many compounds that contain H when dissolved in water are also acids
e.g. H2SO4
H2SO4(aq)
HCl
HCl(aq)
Recall: physical states (s) (l) (g) (aq)
hydrogen sulfate
sulfuric acid
hydrogen chloride
hydrochloric acid

50. Other Acids? CH3COOH(aq) or HC2H3O2 (aq) H2CO3(aq) HCl(aq) HNO3(aq)
Some common ones that you should know
acetic acid
carbonic acid
hydrochloric acid
nitric acid
sulfuric acid
CH3COOH(aq)
or HC2H3O2 (aq)
H2CO3(aq)
HCl(aq)
HNO3(aq)
H2SO4(aq)