Ionic compounds

Ionic Compounds

An ionic compound forms when a metal transfers one or more electrons to a nonmetal. The metal becomes a positively charged cation and the nonmetal becomes a negatively charged anion. The two oppositely charged ions attract each other electrostatically. This electrostatic attraction is the ionic bond. The ratio of cations to anions in the compound is always whatever is needed to make the overall charge of the compound equal to zero. This principle of electrical neutrality is the foundation for writing all ionic formulas.

Figure 1: Formation of an Ionic Bond

Illustrating the formation of an ionic bond, this diagram shows the electron transfer from a sodium atom to a chlorine atom, resulting in the creation of oppositely charged ions that attract each other.

The Charge Neutrality Rule

Every ionic compound is electrically neutral as a whole. The total positive charge from all cations must exactly equal the total negative charge from all anions. This gives us a simple rule for writing formulas:

Charge neutrality condition:

Total positive charge+Total negative charge=0\text{Total positive charge} + \text{Total negative charge} = 0

Writing Formulas for Ionic Compounds

Simple (Monoatomic) Ions

For compounds with simple monoatomic ions, the formula is found by determining the ratio of ions that produces zero overall charge. The most efficient method is the cross-charge method. Use the magnitude of the cation's charge as the subscript of the anion, and the magnitude of the anion's charge as the subscript of the cation. Then reduce to the lowest whole-number terms if necessary.

Table 1. Formula Writing for Simple Ionic Compounds Using the Charge Neutrality Rule

Table displaying the charge balance and chemical formulas of common ionic compounds, including their cation, anion, and corresponding names.

Worked Example

Determine the formula for calcium bromide.

Step 1: Identify the ion charges.

Ca2+\text{Ca}^{2+}

Calcium forms a +2 ion.

Br−\text{Br}^{-}

Bromide forms a −1 ion.

Step 2: Determine the ratio needed for charge neutrality.

One calcium ion contributes +2.

Two bromide ions contribute −2.

Step 3: Verify the charge balance.

1(+2)+2(−1)=01(+2) + 2(-1) = 0

Therefore, the formula is:

CaBr2\text{CaBr}_2

Compounds Containing Polyatomic Ions

Polyatomic ions are treated exactly like monoatomic ions when building compounds. The total charge must still equal zero.

The key rule when writing formulas is that if more than one of the same polyatomic ion is needed, enclose the entire ion in parentheses before writing the subscript. The subscript then applies to the whole group, not just the last element.

Worked Example: Aluminum Sulfate

Al³⁺ combines with SO₄²⁻.

Step 1: Determine the least common multiple of the ion charges.

The least common multiple of 3 and 2 is 6.

Step 2: Determine the number of each ion required.

For aluminum:

6÷3=26 \div 3 = 2

For sulfate:

6÷2=36 \div 2 = 3

Step 3: Verify charge neutrality.

2(+3)+3(−2)=02(+3) + 3(-2) = 0

Therefore, the formula is:

Al2(SO4)3\text{Al}_2(\text{SO}_4)_3

Notice the parentheses around SO₄. Without them, the formula Al₂SO₄₃ would be ambiguous or incorrect. The parentheses make it clear that the subscript 3 applies to the entire sulfate group, meaning the compound contains:

• 2 aluminum atoms

• 3 sulfur atoms

• 12 oxygen atoms

Figure 2: Using Parentheses with Polyatomic Ions

Understanding the importance of using parentheses with polyatomic ions, this diagram contrasts the correct Al₂(SO₄)₃ structure with an incorrect version, highlighting key differences in chemical composition and charge balance.

Table 2. Formulas for Ionic Compounds Containing Polyatomic Ions

This table lists various cations and anions along with their chemical formulas and notes, providing a reference for understanding their compositions and applications.

Properties of Ionic Compounds

The strong electrostatic forces between cations and anions in an ionic lattice give ionic compounds a characteristic set of physical properties that distinguish them from covalent compounds.

High Melting and Boiling Points

Ionic bonds are strong. It takes a large amount of energy to separate ions from the crystal lattice.

Examples:

• NaCl melts at 801°C

• MgO melts above 2800°C

Figure 3: Ionic Crystal Lattice

An illustration of a sodium chloride (NaCl) ionic crystal lattice, demonstrating how cations and anions form a stable, repeating 3D pattern with strong electrostatic attractions, leading to properties like high melting points and brittleness.

Hard but Brittle

The lattice structure is rigid, making ionic solids hard. However, when struck, layers of ions can shift so that ions with the same charge become adjacent to one another. These like charges repel strongly, causing the crystal to fracture.

Figure 4: Brittleness of Ionic Solids

Illustration of ionic solids' brittleness, showing how stress causes layers of ions to shift, leading to repulsion between like charges and resulting in cracks or breaks in the crystal structure.

Solid at Room Temperature

The strong lattice forces hold ions in fixed positions. As a result, ionic compounds generally have definite shapes and definite volumes at room temperature.

Nonconductors in the Solid State

In a solid ionic crystal, the ions are locked into place and cannot move freely. Since charged particles cannot move, electrical current cannot be conducted.

Good Conductors When Melted or Dissolved

When an ionic compound melts or dissolves in water, the ions become free to move. These mobile ions can carry electrical charge through the liquid.

This is why dissolved salts are classified as electrolytes.

Figure 5: Electrical Conductivity of Ionic Compounds

The image explains the electrical conductivity of ionic compounds, highlighting that they conduct electricity when ions are free to move in molten or aqueous states, but not in solid form.

Often Soluble in Water

Water molecules interact strongly with ions through ion-dipole forces, allowing many ionic compounds to dissolve.

However, solubility varies widely. Not all ionic compounds are water-soluble.

Figure 6: Hydration of Ions

Explaining the hydration of ions, this diagram illustrates how water molecules dissolve ionic compounds, surrounding and stabilizing individual ions through ion-dipole attractions, preventing them from rejoining.

Summary

Ionic compounds form when electrons are transferred from a metal to a nonmetal, producing oppositely charged ions that attract through electrostatic forces. The formulas of ionic compounds are governed by the charge neutrality rule, which requires that the total positive and negative charges balance to zero. Formula writing requires determining the proper ratio of ions, whether they are monoatomic or polyatomic. The strong ionic lattice gives ionic compounds their characteristic properties, including high melting points, brittleness, electrical conductivity when molten or dissolved, and frequent solubility in water.

Key Points

• Ionic compounds form through electron transfer between metals and nonmetals.

• Cations are positively charged ions, while anions are negatively charged ions.

• Every ionic compound must obey the charge neutrality rule.

• Cross-charge methods help determine ionic formulas.

• Parentheses are required when multiple polyatomic ions are present.

• Ionic compounds possess strong electrostatic attractions within crystal lattices.

• Ionic compounds typically have high melting and boiling points.

• Ionic solids are hard but brittle.

• Ionic compounds conduct electricity when molten or dissolved, but not as solids.

• Many ionic compounds dissolve in water because of ion-dipole interactions.