Chapter 7
Ionic Compounds and Metals
Section 7.1
Ion formation

Chemical Bonds
A chemical bond is the force that holds two atoms together.
• Can form by the attraction between the
positive nucleus of one atom and the
negative electrons of another
• Can form between positive and negative ions

Valence Electrons
• Electrons in the outermost principal energy level
• Shown in the electron dot structures
• Octet rule – atoms will gain, lose or share electrons to obtain 8 valence electrons
• The valence electrons determine the bonding
properties of the atom

Positive Ion Formation
• A positively charged ion is called a cation.
• Positive ions are formed when an atom loses one or more valence electrons
• Metals make positive ions

Negative Ion Formation
• A negatively charged ion is called an anion.
• Negative ions are formed when an atom gains one or more electrons in its valence shell.
• Nonmetals make negative ions.

7.2: Ionic bonds and ionic compounds
Formation of an Ionic Bond
• An ionic bond is the electrostatic force that holds oppositely charged particles together in an ionic compound
• Compounds that contain ionic bonds are called
ionic compounds.
• Ionic compounds are formed between
metals (+ charge) and nonmetals (- charge).

Binary Ionic Compounds
• Contain a metallic cation and a nonmetallic
anion.
• Formation of Binary Ionic Compounds
– Electron(s) is/are transferred from metal to nonmetal
– Metal becomes positive, nonmetal becomes negative
– Opposite charges attract

Properties of Ionic Compounds
• Take the structure of a crystal lattice
– Many units of positive and negative ions stick together in a three-dimensional geometric arrangement
• Can conduct electricity when dissolved in water
(they are electrolytes and break into ions when dissolved in water), but not in solid form
• Melting point, boiling point and hardness depend upon how strongly the ions are attracted to each other

Formulas for Ionic Compounds
• Monatomic ions are one-atom ions
– Examples: Mg 2+ , Br -1
• Oxidation numbers are the charges on ions
– Note: some elements have multiple oxidation states – you will have a periodic table to tell this
• Binary ionic compounds are made of two monatomic ions (one positive, one negative)

Formulas for Binary Ionic
Compounds
• Symbol for cation is written first, anion second
• Subscripts tell the number of atoms of each element
• What are the following compounds made of?
– CaF
2
1 calcium, 2 fluorine
– Na
2
S 2 sodium, 1 sulfur
– NaCl 1 sodium, 1 chlorine

Naming Binary Ionic Compounds
• Name the cation first
• Name the anion second with –ide at the end
• Examples
– CaF
2
 calcium fluoride
– Na
2
S  sodium sulfide
– NaCl  sodium chloride

• K
2
O
Try Naming a few more
Binary Ionic Compounds potassium oxide
• Al
2
S
3
• Na
3
N aluminum sulfide sodium nitride

What if the cation has more than one oxidation state?
• You tell which ion was used by putting a Roman Numeral after the name of the cation
• Example:
– CuS
• We know S was -2 (that’s the only one it makes)
• If there is only one atom of each element, the Cu must have been +2
• So, the name is written as Copper (II) sulfide [the “II” indicates the charge]
• Make sure, especially with transition elements, that you are checking the oxidation states

Writing Formulas for Binary
Ionic Compounds
• Look up the charges for each element
• For a compound to form, the total charge must balance out to zero (positive charges must equal negative charges)
• Example:
– Sodium bromide
• Na is +1, Br is -1
• Only need one of each to balance
• Formula is NaBr

Try writing some more formulas
Binary Ionic Compounds
• Potassium Iodide KI
• Aluminum bromide
• Magnesium chloride
• Cesium nitride
AlBr
3
MgCl
2
Cs
3
N

Formulas for Polyatomic
Ionic Compounds
• Polyatomic ions are ions that are made up of more than one atom
• You will have a chart for these and do not have to memorize them.
• Examples:
– SO
4
2= sulfate
– CN = cyanide
– NH
4
+ = ammonium

Naming Polyatomic
Ionic Compounds
• Name the cation first, anion second
• Name the polyatomic as is – don’t change its name at all
• Examples:
– Ca
3
(PO
4
)
2
– Mg(CN)
2
– NH
4
Cl calcium phosphate magnesium cyanide ammonium chloride

Now you try naming
Polyatomic Ionic Compounds
• NaNO
3 sodium nitrate
• Ca(ClO
3
)
2
• Al
2
(CO
3
)
3 calcium chlorate aluminum carbonate

Writing formulas for
Polyatomic Ionic Compounds
• Same as binary ionic compounds EXCEPT you may not change anything in the polyatomic ion formula
• Put them in a (parenthesis) and put subscripts outside that parenthesis
• Example:
– Calcium Nitrate
• Ions are Ca 2+ and NO
3
-
• Formula will be Ca(NO
3
)
2

Now you try writing formulas for
Polyatomic Ionic Compounds
• Sodium hydroxide NaOH
• Copper (II) nitrate
• Silver chromate
Cu(NO
3
)
2
Ag
2
CrO
4

7.3: Metallic bonds and the properties of metals
• The electron sea model proposes that all the metal atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons
• Since the electrons are free to move, they are called delocalized electrons
• A metallic bond is the attraction of a metallic cation for delocalized electrons

Properties of Metals
(revisited)
• Moderately high melting points
• High boiling points
• Malleable, ductile, durable
• Conduct heat and electricity well
• Transition metals are harder/stronger than alkali metals because the transition metals have more delocalized electrons