Worksheet 6.3, 6.4 - Ions and Ionic Compounds

For group A representative elements, the group number gives the number of electrons in the outermost energy level. These electrons are the ones that participate in chemical bonding and are called valency electrons.

1. How many valency electrons do each of the following elements have?

a) potassium b) sulfur c) nitrogen

d) calcium e) iodine f) neon

An ion is a charged atom. If an atom loses electrons, it will have more protons than electrons and therefore have a positive charge forming a cation. If an atom gains electrons, it will have more electrons than protons and therefore have a negative charge forming an anion.

Remember that metals tend to lose electrons and nonmetals to gain electrons to achieve the octet or 8 electrons in the outer energy level of a noble gas because this is very stable. Metals will then lose their valency electrons forming cations and nonmetals will gain enough electrons to complete an octet forming anions.

Examples: Write the chemical shorthand (chemical symbol with the charge as a superscript) for the most common ions of the following elements. The charge on an ion refers to the number of electrons lost or gained.

sodium – Group 1A, loses 1 valency electron so Na⁺ (a “1” is understood in front of the +).

chlorine– Group 7A, has 7 valency electrons, gains 1 electron to make 8 so Cl^– (a “1” is understood in front of the –)

magnesium– Group 2A, loses 2 valency electrons so Mg²⁺

phosphorus– Group 5A, has 5 valency electrons, gains 3 electrons to make 8 so P^3–

2. Write the chemical shorthand for the ions of

a) potassium b) sulfur c) nitrogen

d) calcium e) iodine

3. Would neon normally form ions? Why or why not?

Formulas of ionic compounds:

The atoms in molecules bond to one another through "sharing" of electrons. Ionic compounds on the other had have atoms or molecules that bond to one another through their mutual attraction between positive and negative charges:

Positively charged cations and negatively charged anions attract one another. Conversely, cations repel one another as do anions.

Electrostatic attraction is indiscriminate. That is a cation can attract more than one anion and visa versa. The result is that cation-anions attractions form a large array that we call an ionic compound or "salt". The bonds holding these ions together are called "ionic" bonds. However, this array has a very specific composition completely dictated by the charges on the cations and anions.

Formulas of ionic compounds:

The composition of ionic compounds is determined by the requirement that the compounds must be electrically neutral. That is that the charges of the cations and anions must balance or 'cancel" out one another. For example consider sodium cations (Na⁺) and Chlorine anions (Cl^–). Sodium has a positive 1 charge and chloride has a negative 1 charge. Thus one sodium cation cancels one chloride anion (+1 + –1 = 0) resulting in the formula Na₁Cl₁ or NaCl. This formula is called the formula unit since it represents only one unit of the vast NaCl array or lattice.

A small representative bit of a NaCl lattice looks like this:

If you look at the diagram carefully, you will see that the sodium ions and chloride ions alternate with each other in each of the three dimensions.

This diagram is easy enough to draw with a computer, but extremely difficult to draw convincingly by hand. We normally draw an “exploded” version which looks like this:

Naming an ionic compound:

In the name of an ionic compound, the cation is listed first and the anion, second. Anionsr usually have an –ide suffix. Therefore NaCl is named sodium chloride.

Example: Give the formula of the ionic compound, magnesium phosphide.

As shown before, the ions of magnesium and phosphorus are Mg²⁺ and P^3–. To form a neutral formula unit, we would need:

Mg²⁺ P^3–

Mg²⁺ P^3–

Mg²⁺

total cation charge = (3)(+2) = +6 total anion charge (2)(–3) = –6

total charge on one formula unit = (6+) + (6–) = 0

Since we used three Mg²⁺ and two P^3– the formula is Mg₃P₂. The subscripts give the number of each ion used.

Example: Give the name of the ionic compound whose formula is MgO.

Write the name of the cation first and then anion with –ide suffix so magnesium oxide

Transition elements can form ions with several different charges so these must be looked up on a chart (see Table 1 at the end of the handout). For example, iron can be Fe²⁺ and Fe³⁺, called iron (II) ion and iron (III) ion respectively. Note that the Roman numeral following the metal name is the same as the charge of the ion.

Example: Give the formula of the ionic compound, iron (III) chloride.

Fe³⁺ Cl^–

Cl^–

Cl^–

total cation charge = (1)(+3) = +3 total anion charge (3)(–1)= –3

total charge on one formula unit = (+3) + (–3) = 0

Example: Give the name of the ionic compound with the formula CuCl₂.

From the subscripts, the compound contains one copper ion and two chloride ions.

Chlorine is in group 7A, has 7 valency electrons and gains 1 (8-7=1) to form Cl^– ions.

2 Cl^– ions have a charge of (2)( –1) = –2.

1 copper ion must then have a charge of +2 to make a neutral compound, (–2) + (+2) = 0

The compound name is copper (II) chloride. The Roman number in parentheses is the charge on the copper ion.

Below are other examples of simple ionic compounds made up of single atom (monatomic) ions:

Polyatomic ions and compounds:

The ionic compounds above were made up of monatomic ions. Molecules can also be ions - polyatomic ions. Most polyatomic ions are anions with one notable exception - the ammoniun cation (NH⁴⁺). The composition of salts with polyatomic ions is determined by the same rule as with monatomic ions. A few example are given below:

Note that when there is two or more polyatomic ions in the formula the ion is enclosed with parenthesis "()".

Formulas and names of ionic compounds

Summary

• Positively charged ions are called cations
• Negatively charged ions are called anions
• The valency electrons of Group A elements is given by the group number.
• Group A metals lose their valency electrons so the charge of the ion is + and equal to the group #. For example, beryllium is in Group 2A and the ion is Be²⁺.
• Group A nonmetals gain electrons to add to their valency electrons to gain an octet. The charge of the ion is – and equal to (8 – group #). For example, sulfur is in Group 6A, 8-6=2 so the ion is S^2–.
• The formula of an ionic compound represents the simplest whole number ratio of ions present which is one formula unit
• The net charge on an ionic compound is zero, so the sum of the positive charges equals the sum of the negative charges.
• A subscript number written to the right of an element's symbol tells us how many of those ions are present in the formula. For example: potassium sulfide has the formula K₂S, 2 potassium ions (K⁺) and 1 sulfide ion (S^2–) are present.
• If no number is given, then only one of that ion is present. For example: sodium chloride has the formula NaCl, 1 sodium ion (Na⁺) and 1 chloride ion (Cl^–) are present.
• If more than one polyatomic ion (ion having more that one element in its formula) is present, round brackets enclose the formula of the polyatomic ion and a subscript number written to the right of the final bracket tells us how many of that polyatomic ion are present in the formula. For example, the nitrate ion is a polyatomic ion, NO^3–, it is an ion containing nitrogen and oxygen atoms. Calcium nitrate has the formula Ca(NO₃)₂ and it contains 1 calcium ion (Ca²⁺) and 2 nitrate ions (NO^3–).
• When naming an ionic compound, the name of the cation is given first and the anion second.
• Single element anions have an –ide suffix – example: sodium chloride.
• For cations of transitions elements, the charge is written as a roman numeral in parentheses after the name of the cation – example iron (III) chloride.

4. Write the formulas for the following ionic compounds.

name of compound	cation	anion	formula
sodium sulfide
calcium iodide
beryllium phosphide
aluminum chloride
cobalt (II) oxide
lead (IV) sulfide
lithium floride
tin (IV) chloride
manganese (III) oxide
lead (II) sulfide

5. Write the formulas for the following ionic compounds containing polyatomic ions.

name of compound	cation	anion	formula
lithium carbonate
calcium nitrate
ammonium sulfide
sodium sulfate
sodium thiosulfate
mercury (II) chlorate
potassium dichromate
tin (IV) phosphate
lead (II) nitrite
mercury (I) sulfide

6. Write the names for the following ionic compounds given the formula.

formula	cation	anion	name of compound
NH4Cl
Fe(NO3)3
TiBr3
Cu3P
SnSe2
GaAs
Pb(SO4)2
Be(HCO3)2
Mn2(SO3)3
Al(CN)3

Table 1: Metal cations with more than one common charged form

Fe^2+	iron(II) ion
Fe^3+	iron(III) ion
Cu^+	copper(I) ion
Cu^2+	copper(II) ion
Hg2^2+	mercury(I) ion
Hg^2+	mercury(II) ion
Pb^2+	lead(II) ion
Pb^4+	lead(IV) ion
Sn^2+	tin(II) ion
Sn^4+	tin(IV) ion
Cr^2+	chromium (II) ion
Cr^3+	chromium (III) ion
Mn^2+	manganese (II) ion
Mn^3+	manganese (III) ion
Co^2+	cobalt (II) ion
Co^3+	cobalt (III) ion

Table 2: Common polyatomic cations, arranged by family. (alternate names in italics)

carbon

nitrogen

sulfur

chlorine

CO3^2^–	carbonate
HCO3^–	hydrogen carbonate (bicarbonate)

NO3^–	nitrate
NO2^–	nitrite

SO4^2^–	sulfate
SO3^2^–	sulfite
S2O3^2^–	thiosulfate
HSO4^–	hydrogen sulfate (bisulfate)
HSO3^–	hydrogen sulfite (bisulfite)

ClO4^–	perchlorate
ClO3^–	chlorate
ClO2^–	chlorite
ClO^–	hypochlorite

phosphorus

cyanide

cations

metal oxyanions

PO4^3^–	phosphate
HPO4^2^–	hydrogen phosphate
H2PO4^–	dihydrogen phosphate

CN^–	cyanide
OCN^–	cyanate
SCN^–	thiocyanate

NH4^+	ammonium
H3O^+	hydronium
Hg2^2+	mercury(I)

CrO4^2	chromate
Cr2O7^2^–	dichromate
MnO4^–	permanganate

oxygen

organics

OH^–	hydroxide
O2^2^–	peroxide

C2H3O2^–

acetate