The Formation of Ionic and Covalent Bonds (2.1)
Clues in Naturally Occurring Compounds
▪ ores are mined to extract pure metals (metal ion and a non-metal ion bonded together)
▪ most metals are found in ores, precious metals (ie. gold and silver) are found in elemental form
▪ based on observing the atmosphere we see that: metals usually form bonds with non-metals (make solids),
non-metals can bond with non-metals (same or different) in any state, noble gases are the only elements
NEVER found in a combined form in nature
Stability of Atoms and the Octet Rule
▪ full valence shell = stable (fig 2.3)
▪ when bonds form between atoms, they do so in a way that gives each atom a filled valence shell (called
octet)
▪ octet rule: when bonds form between atoms, the atoms gain, lose or share electrons in such a way that they
create a filled outer shell containing eight electrons
▪ transition elements are exceptions to the rule (more than 8), hydrogen only has 2
The Formation of Ionic Bonds
▪ ionic bond: between metal (+ve) and non-metal (-ve) ions (electrostatic force), forming an ionic compound
▪ the total number of electrons lost = the total number of electrons gained
Example:
**Remember the criss-cross rule?**
Ionic Compounds Containing Transition Metals
▪ the number of electrons that can be lost can change, see the periodic table (first one is most common)
▪ for questions you will always be given enough information to know/figure out which one to use
▪ Lewis dot diagrams for transition elements only show the electrons that can be lost
Example:
The Formation of Covalent Bonds
▪ shared electrons, molecular compound, non-metals only
▪ only unpaired electrons participate in bonds
▪ can have more than 2 elements (ie. CO2)
Example:
Types of Covalent Bonds and Electron Pairs
▪ single bond: one pair of electrons
▪ double bond: two pairs of electrons
▪ triple bond: three pairs of electrons
▪ bonding pair: electrons that are shared
▪ lone pair: electrons that are not involved in the bond
Examples:
**these are called Lewis Structures**
Drawing Lewis Structures
Step 1: Identify central atom (highest bonding capacity), put all other symbols around it
Step 2: Add up valence electrons
Step 3: Place one pair of electrons between each pair of
atoms
Step 4: Place remaining electrons as lone pairs on atoms
(not central)
Step 5: Place remaining electrons on central atom, in pairs
Step 6: If the central atom does not have a full octet, move
lone pairs from the surrounding atoms into a bonding
position with the central atom (to make double or triple
bonds)
Polyatomic Ions and Bond Formation
▪ polyatomic ion: molecular compound with an excess or
deficit of electrons (ie. CO32-)
▪ a correct polyatomic ion drawing had square brackets and
charge (fig 2.13B)
▪ form ionic bonds (fig 2.14)
Electronegativity Difference and Bond Type
▪ electrons generally spend more time around the atoms with greater electronegativity
Example:
▪ covalent bonds with unequal electron distribution are polar covalent
▪ bond type is determined by electronegativity difference
Electronegativity Difference
▪ subtract one from the other (see periodic table)
▪ > 1.7 = mostly ionic
▪ 0.4-1.7 = polar covalent
▪ <0.4 = slightly polar covalent
▪ 0 = non-polar covalent (equal sharing)
Example:
Percent Ionic and Covalent Character
▪ different way of classifying bonds
▪ Table 2.1 *you DO NOT need to memorize this*
Classwork:
Page 63 #3, 4, 7, 10, 12