Chapter 5
Ions and Ionic Compounds
Unit Essential Question:
What are the characteristics of
ionic compounds?
Lesson Essential Question:
Why do atoms form ions?
Section 1: Simple ions
 How reactive atoms are depends upon the
number of valence electrons.
 Certain number makes atoms stable.
 Octet rule – elements with full outer energy
levels tend to be stable. Full energy levels
have eight electrons.
 All noble gases have a satisfied octet rule
except He (2 valence electrons).
 Don’t want to gain or lose an electron.
Stability
 This will be an important theme in this chapter
and in the next.
 Being stable is the goal of atoms and is the
reason why bonding takes place.
 By bonding, atoms change their electron
configurations to become more stable.
 There are several terms/phrases that you will see
that indicate an atom is stable:
 #1: Octet rule is satisfied/achieved.
 #2: The atom has a noble gas configuration.
 #3: The atom has a full valence/outer shell.
Alkali Metals & Halogens
 Most reactive elements.
 Alkali metals-1 valence e- ; halogens- 7 valence e Full valence shell is achieved if alkali metals lose one e-
and halogens gain one e-.
 Often bond together- the alkali metal gives an e- to the
halogen.
 Example:




K: [Ar]4s1 and Cl: [Ne]3s23p5
K will lose one e- and Cl will gain one e-.
Each now has the same configuration: [Ne]3s23p6
Both have the same configuration as the noble gas
argon- they’re now stable, have an octet!
Ions
 Ion – atom or molecule that has a charge because it has
gained or lost electrons. Identity does not change!
 Cation – ion with a positive charge; e- are lost.
 The ‘t’ in cation resembles a + sign. This can help you remember
that cations are positive.
 Anion – ion with a negative charge; e- are gained.
 Ions form as atoms change their valence e- to try and
achieve a noble gas configuration.
 As we saw before:
 K = [Ar]4s1  loses 1 e- ; K is now a K+ ion.
 Cl = [Ne]3s23p5  gains 1 e- ; Cl is now a Cl- ion.
Patterns of Charges on the Periodic Table
 Since elements
in the same group
have the same #
of valence e-,
elements in the
same group often
form the same
ions (in the s and
p blocks).
+1 +2
+3
-3
-2
-1
*Metals tend to lose e- (become cations) and
nonmetals tend to gain e- (become anions).
No patterns in the d-block
 Transition metals do not follow a pattern according to
group number.
 Almost all transition metals form more than one cation.
Pb4+
Ions- what changes and what doesn’t
Ions and Ionic compounds
 Properties change!
 Compounds formed from ions have very different
properties from the original atoms.
 Ex: sodium and chlorine vs. sodium chloride.
Remember:
 Identities don’t change!
 Ions do not BECOME noble gases!
 Have electron configurations of noble gases and
therefore behave like noble gases.
 They become more stable.
Lesson Essential Question:
What properties result from
forming ionic bonds?
Section 2: Ionic Bonding and Salts
 Ionic bonds form when oppositely charged ions attract.
 Salt – name given to many ionic compounds.
 Ex: NaF, KCl, CaSO4, etc.
 Thousands of ionic compounds are called salts.
 Ionic compounds as a whole are neutral in charge.
 Not just one cation and one anion come together- many
are attracted.
 Results in a tightly packed crystal lattice
structure.
Energy and Ionic Bonding
 There are numerous steps to forming ionic
compounds from elements (ie: forming NaCl from
Na and Cl).
 Energy is involved in every step.
 Overall, more energy is released than is
absorbed.
 This makes forming ionic compounds
spontaneous (favorable).
 Lattice energy: energy released when ionic
bonds form.
Properties of Ionic Compounds
 Attractive forces (+ and -) make bonds very strong.
 High melting and boiling points.
 Usually solids at room temperature.
 Conduct electricity when melted (molten) or in solution
(dissolved in water).
 Most dissolve well in water.
 Hard and brittle as crystals.
 Due to repeating pattern of ions.
 Crystal lattice – regular pattern in which a salt crystal
is arranged.
 Made up of many repeating units (smallest called a unit
cell).
Lesson Essential Question:
How are ionic bonds represented
using formulas and names?
Section 3: Names and Formulas
of Ionic Compounds
Naming Monatomic Ions
 Monatomic ions – one atom with a charge.
 Mono = one.
 Ex: Al+3, O-2
 Cation names do not change. Add ‘ion’ after the name.
 Na+ = sodium ion; Ba+2 = barium ion
 Transition metals can form more than one ion. In order
to tell them apart Roman Numerals are used.
Value of Roman
 Cu+ = copper(I); Cu+2 = copper(II)
Numeral tells
+2
+4
 Sn = tin(II); Sn = tin(IV)
you the size of
the charge.
 Anion names do change.
 “-ide” ending is added at the end of the element name.
 S-2 = sulfide; N-3 = nitride; F- = fluoride
Naming Polyatomic Ions
 Polyatomic ions – 2 or more atoms with a charge.
 Poly = many.
 Ex: NH4+, SO4-2 (ammonium, sulfate)
 These are names that you must memorize.
Ion Chart
 You should know all cations and all starred* anions on the
ion chart.
 Remember- monatomic ions in the s and p blocks can be
determined just by looking at the periodic table! These are
not included on the ion chart but you are expected to know
them!
 There will be a quiz on ion names and charges!
Naming Ionic Compounds
 ALWAYS name and write the cation before the anion.
 Binary – compounds with 2 elements.
 Use rules for naming monatomic ions.
 Ex: NaCl: sodium chloride (Na+ and Cl-)
CaBr2 calcium bromide (Ca+2 and Br-)
 Recall use of Roman numerals for transition metal
cations with multiple possible charges!
 Ex: FeCl3: iron (III) chloride (Fe+3 and Cl-)
Cu2O: copper (I) oxide (Cu+ and O-2)
 To help you figure out the charge:

For right now looking at the subscript for the anion will
help you determine the charge on the transition metal
cation.
Naming ionic compounds cont.
 Naming compounds with 3 or more elements:
 Use rules for naming monatomic ions AND polyatomic ion
names.
 Ex: K3PO4: potassium phosphate (K+ and PO4-3)
(NH4)2S : ammonium sulfide (NH4+ and S-2)
 Notice that parentheses are placed around polyatomic ions
whenever there are more than one used in the chemical
formula.
o Otherwise looks like: NH42S
 Another ex: NH4NO3 : ammonium nitrate (NH4+ and NO3-)
Writing Chemical Formulas for Ionic
Compounds
 Compounds must be neutral!
aluminum oxide
 Positive charge = negative charge.
 Use the “criss-cross” method:
 Step 1: determine ions from the name.
Step 1: Al+3
O-2
Step 2: Al+3
 Step 2: criss-cross numbers only.
O -2
Step 3: Al2O3
 Step 3: be sure to give the lowest ratio possible.
 Reduce if necessary! (ex: Sn2O4  SnO2)
 Note that the positive ion is written first!
 Use parentheses if there is more than one polyatomic ion!
 Ex: calcium nitrate: Ca(NO3)2 (NOT: CaNO32)
 Don’t write 1’s in formulas.
 We assume there’s 1 if there’s no number written after the
charge.