Electron transfer, cations and anions, charge balance, crystal lattices, formula units, properties of ionic compounds, electrolytes, and the role of ionic bonding in minerals, biology, and technology
Table salt is one of the simplest and most familiar substances on Earth. People have mined it, traded it, preserved food with it, and dissolved it in water for thousands of years. Yet if you look at salt through the lens of chemistry, you find something remarkable: table salt is not made of salt molecules. It does not contain neutral sodium atoms or neutral chlorine atoms. Instead, it is a three-dimensional arrangement of charged particles — sodium ions with a positive charge and chloride ions with a negative charge — locked together by the powerful electrical attraction between opposite charges.
That attraction is ionic bonding.
Ionic bonding is one of the fundamental ways atoms combine to form chemical compounds. It occurs when electrons are transferred from one atom to another, creating oppositely charged ions that attract each other strongly. The result is a class of compounds with distinctive properties: high melting points, crystalline solid structure, the ability to conduct electricity when dissolved in water, and characteristic brittleness when struck.
In Chapter 7 you learned that electrons — especially valence electrons — determine how atoms behave chemically. In Chapter 8 you learned that the periodic table organizes elements by their electron arrangements and that metals and nonmetals behave very differently. Chapter 9 brings those ideas together. You will learn why metals tend to lose electrons, why nonmetals tend to gain them, how the transfer of electrons creates ions, and how those ions organize themselves into the ionic compounds that form minerals, salts, ceramics, and countless materials essential to life and technology.
By the end of this chapter you will be able to explain how ionic bonds form, write correct chemical formulas for ionic compounds, predict the charges of common ions, describe the structure and properties of ionic compounds, and connect what you know about ions to real materials in the world around you.
Sodium is a soft, silver metal that reacts violently with water. Chlorine is a toxic yellow-green gas. Neither substance is safe to handle carelessly. Yet when these two elements react, the result is sodium chloride — table salt — which is safe, stable, and essential for human life. How can two dangerous elements combine to produce something harmless and useful? By the end of this chapter you will understand exactly why, at the level of electrons and charges.
Every atom contains a nucleus with protons and neutrons surrounded by electrons in energy levels. The electrons in the outermost energy level — the valence electrons — are the ones that participate in chemical bonding. Noble gases in Group 18 have full outer electron shells and are chemically unreactive. Chemists describe this arrangement as stable.
Other elements are not naturally stable in this way. Sodium has one valence electron. Chlorine has seven. Neither arrangement is as stable as the eight valence electrons that characterize the noble gas argon. This tendency — atoms behaving in ways that give them eight valence electrons in their outer shell — is called the octet rule. It is a useful guide that explains an enormous amount of chemical behavior.
Atoms tend to gain, lose, or share electrons in order to achieve a full outer shell of eight valence electrons — the stable arrangement found in noble gases. Atoms that cannot easily reach an octet by sharing electrons may instead transfer electrons completely, producing ionic bonding.
When a metal atom gives up one or more valence electrons to a nonmetal atom, both atoms are transformed. The metal atom, having lost electrons, now has more protons than electrons. It carries a net positive charge. A positively charged ion is called a cation.
The nonmetal atom, having gained extra electrons, now has more electrons than protons. It carries a net negative charge. A negatively charged ion is called an anion.
Once a cation and an anion exist near each other, opposite charges attract. This electrostatic attraction — the force between opposite charges — is the ionic bond.
It is important to understand that an ionic bond is not a molecule in the usual sense. In an ionic compound, ions are arranged in a repeating pattern in three dimensions, with each cation surrounded by several anions and each anion surrounded by several cations. The compound has no individual molecules, only a continuous network of ions held together by electrostatic attraction.
| Particle | Charge | How It Forms | Example |
|---|---|---|---|
| Cation | Positive (+) | Atom loses one or more electrons | Na → Na⁺ + e⁻ |
| Anion | Negative (−) | Atom gains one or more electrons | Cl + e⁻ → Cl⁻ |
| Ionic bond | — | Electrostatic attraction between cation and anion | Na⁺ + Cl⁻ → NaCl |
Sodium is a reactive metal. Chlorine is a toxic gas. When they react, electrons transfer and ionic bonding forms sodium chloride — table salt — which is safe and stable. The properties of an ionic compound are completely different from the properties of the elements that formed it. Ionic bonding does not just rearrange atoms; it creates an entirely new substance.
Metals occupy the left side and central block of the periodic table. Most metals have one, two, or three valence electrons. When metal atoms react chemically, they tend to lose these valence electrons. Losing electrons reduces the electron count while leaving the proton count unchanged, producing an ion with a positive charge equal to the number of electrons lost.
Na → Na⁺ + e⁻ (charge = +1)
Mg → Mg²⁺ + 2e⁻ (charge = +2)
Al → Al³⁺ + 3e⁻ (charge = +3)
Nonmetals are found in Groups 14 through 17. They have many valence electrons — typically five, six, or seven — and are much closer to a full outer shell than metals are. When nonmetal atoms react with metals, they gain electrons to complete their outer shell, producing anions with a negative charge equal to the number of electrons gained.
Cl + e⁻ → Cl⁻ (charge = −1)
O + 2e⁻ → O²⁻ (charge = −2)
N + 3e⁻ → N³⁻ (charge = −3)
| Group | Electrons Lost or Gained | Ion Charge | Examples |
|---|---|---|---|
| Group 1 (metals) | Lose 1 | +1 | Li⁺, Na⁺, K⁺ |
| Group 2 (metals) | Lose 2 | +2 | Mg²⁺, Ca²⁺, Ba²⁺ |
| Group 13 (metals) | Lose 3 | +3 | Al³⁺ |
| Group 16 (nonmetals) | Gain 2 | −2 | O²⁻, S²⁻ |
| Group 17 (halogens) | Gain 1 | −1 | F⁻, Cl⁻, Br⁻, I⁻ |
| Group 18 (noble gases) | Rarely form ions | — | He, Ne, Ar |
Transition metals in the central block of the periodic table can often form more than one possible ion. Iron forms both Fe²⁺ and Fe³⁺. Copper forms Cu⁺ and Cu²⁺. Because of this variability, the charges of transition metal ions must often be stated explicitly or determined from context. You will learn how to handle transition metals in chemical naming in Chapter 11.
Cations take the name of the element followed by "ion." Sodium forms the sodium ion (Na⁺). Magnesium forms the magnesium ion (Mg²⁺).
Anions take the root of the element name followed by the suffix "-ide," plus "ion." Chlorine forms the chloride ion (Cl⁻). Oxygen forms the oxide ion (O²⁻). Sulfur forms the sulfide ion (S²⁻). Nitrogen forms the nitride ion (N³⁻). Fluorine forms the fluoride ion (F⁻). Bromine forms the bromide ion (Br⁻).
An ionic compound as a whole must be electrically neutral. The total positive charge from all cations must exactly equal the total negative charge from all anions. This is not a convention — it reflects the physical reality that ordinary chemical compounds carry no net electrical charge.
To write a correct ionic compound formula: (1) write the cation and its charge, then the anion and its charge; (2) find the ratio of ions that produces zero net charge; (3) use that ratio as the subscripts in the formula.
1 × (+1) + 1 × (−1) = 0 ✓ → ratio is 1:1
One Mg²⁺ = +2
Two Cl⁻ = −1 + (−1) = −2
Total: +2 + (−2) = 0 ✓
Two Al³⁺ = +6
Three O²⁻ = −6
Total: +6 + (−6) = 0 ✓ → ratio is 2:3
A helpful technique: write the ions and their charges, then use each ion's charge number (without the sign) as the subscript for the other ion. If the resulting subscripts share a common factor, reduce them to the lowest ratio. Always reduce — Ca₂O₂ must become CaO.
Charge of Al = 3 → O gets subscript 3
Charge of O = 2 → Al gets subscript 2
Result: Al₂O₃ (already lowest terms)
Crossover gives Ca₂O₂
Both subscripts share factor 2 → reduce to CaO
Final formula: CaO
If the crossover method gives you Ca₂O₂, you must reduce this to CaO. The formula must represent the simplest whole-number ratio of ions. Writing Ca₂O₂ instead of CaO implies a different structure and is chemically incorrect.
Write the correct ionic compound formula for each pair. Show the charge-balancing step.
Ionic compounds do not exist as individual molecules. Chemists use the term formula unit to describe the simplest ratio of ions in the compound. The formula unit of NaCl tells you that for every sodium ion there is one chloride ion. The formula unit of MgCl₂ tells you that for every magnesium ion there are two chloride ions.
The formula unit is a ratio, not a discrete particle. In a crystal of table salt, there are no individual NaCl units. There is instead a massive, three-dimensional arrangement of Na⁺ and Cl⁻ ions in a 1:1 ratio extending in all directions.
| Formula Unit | Ions Present | Ratio of Ions |
|---|---|---|
NaCl | Na⁺ and Cl⁻ | 1 Na⁺ : 1 Cl⁻ |
MgCl₂ | Mg²⁺ and Cl⁻ | 1 Mg²⁺ : 2 Cl⁻ |
Al₂O₃ | Al³⁺ and O²⁻ | 2 Al³⁺ : 3 O²⁻ |
CaO | Ca²⁺ and O²⁻ | 1 Ca²⁺ : 1 O²⁻ |
"Molecule" implies a discrete, bounded particle — a specific number of atoms bonded together. Ionic compounds do not exist as discrete particles. The formula unit is the correct term because it describes only the ratio of ions, not an individual particle. Using "molecule" for an ionic compound like NaCl is chemically inaccurate.
The three-dimensional arrangement of ions in an ionic compound is called the crystal lattice. Every ion in the lattice is surrounded by ions of the opposite charge. This arrangement maximizes the attractive forces between opposite charges and minimizes the repulsive forces between like charges.
In sodium chloride, each Na⁺ ion is surrounded by six Cl⁻ ions — one above, one below, and four to the sides. Each Cl⁻ ion is surrounded by six Na⁺ ions in the same pattern. This regular cubic arrangement extends throughout the entire crystal, giving table salt its characteristic cubic crystal shape. If you look at salt crystals under a magnifying glass, you can often see their cubic form directly.
The lattice energy is the energy released when gaseous ions come together to form one mole of an ionic solid. The more highly charged the ions and the smaller their size, the stronger the attraction, the more energy is released, and the more stable the lattice. Al₂O₃ (with Al³⁺ and O²⁻) has a much higher lattice energy than NaCl (with Na⁺ and Cl⁻), which is why aluminum oxide melts at 2,072°C while sodium chloride melts at 801°C.
When an ionic crystal is struck, layers of ions shift. After the shift, ions of the same charge end up adjacent to each other. The repulsion between like charges causes the crystal to split along flat surfaces. This is why you can cleave an ionic crystal cleanly — and why it shatters rather than bends.
Because all ionic compounds share the same fundamental structure — charged ions in a crystal lattice held by electrostatic attraction — they also share a characteristic set of physical properties. These properties distinguish ionic compounds from other types of substances, particularly from covalent molecular compounds.
| Property | Typical Ionic Compounds | Explanation |
|---|---|---|
| Melting point | Very high (hundreds to over 2,000°C) | Must overcome strong electrostatic attraction between many ions in the lattice |
| State at room temperature | Solid | Melting points far exceed room temperature |
| Crystal structure | Ordered geometric crystals | Ions arrange in precise repeating lattice geometry |
| Brittleness | Hard but shatters rather than bends | Shifting layers align like charges; repulsion fractures the crystal |
| Electrical conductivity | Solid: poor; dissolved or melted: good | Free ions needed to carry charge; solid ions locked in place |
| Solubility | Many dissolve in water; few dissolve in nonpolar solvents | Polar water interacts with ions; nonpolar solvents cannot |
NaCl (ionic): melts at 801°C
CaO (ionic): melts at 2,613°C
Al₂O₃ (ionic): melts at 2,072°C
H₂O (molecular): melts at 0°C
Ethanol (molecular): melts at −114°C
CO₂ (molecular): sublimes at −78.5°C
When you stir table salt into water, the crystals disappear. The sodium and chloride ions have not vanished — they have separated and dispersed throughout the water. This process is called dissociation.
Water molecules are polar: the oxygen end carries a partial negative charge (δ−) and the hydrogen ends carry partial positive charges (δ+). When water contacts an ionic crystal, the negative ends of water molecules attract cations and the positive ends attract anions, gradually pulling ions away from the crystal surface. Once surrounded by water molecules, an ion is said to be hydrated. As more ions are pulled away and hydrated, the crystal dissolves.
Chemists represent dissociation with a chemical equation. The state symbol (s) means solid; (aq) means aqueous — dissolved in water.
NaCl(s) → Na⁺(aq) + Cl⁻(aq)
One formula unit produces 2 ions total.
MgCl₂(s) → Mg²⁺(aq) + 2 Cl⁻(aq)
One formula unit produces 3 ions total.
Ca(NO₃)₂(s) → Ca²⁺(aq) + 2 NO₃⁻(aq)
Subscript 2 in the formula becomes a coefficient 2 in the equation.
Al₂(SO₄)₃(s) → 2 Al³⁺(aq) + 3 SO₄²⁻(aq)
One formula unit produces 5 ions total.
To count the total ions produced per formula unit: add all the coefficients on the right side of the dissociation equation. NaCl produces 2 ions; MgCl₂ produces 3 ions; Al₂(SO₄)₃ produces 5 ions. This matters for calculating how dissolved compounds affect boiling points, freezing points, and electrical conductivity of solutions.
Write the dissociation equation for each compound. Then state the total number of ions produced per formula unit.
When you dissolve table salt in water and connect electrodes to the solution, electricity flows through the liquid. When you dissolve sugar in water and do the same test, no significant current flows. This difference reveals a fundamental distinction between types of dissolved substances.
Electrolyte: a substance that produces ions in solution and therefore allows the solution to conduct electricity. Ionic compounds that dissolve are electrolytes — dissociation releases free ions that carry electrical charge through the solution.
Nonelectrolyte: a substance that dissolves in water without producing ions. Sugar, ethanol, and most molecular compounds that dissolve in water are nonelectrolytes. Their solutions do not conduct electricity because there are no free ions to carry charge.
| Type | Ions in Solution? | Conducts Electricity? | Example |
|---|---|---|---|
| Strong electrolyte | Yes — dissociates completely | Yes, well | NaCl, KBr, MgCl₂ |
| Weak electrolyte | Partially — only some dissociate | Poorly | Acetic acid (vinegar) |
| Nonelectrolyte | No | No | Sugar (C₆H₁₂O₆), ethanol |
A strong electrolyte dissociates completely when dissolved. Every formula unit produces free ions. Ionic compounds that dissolve readily are strong electrolytes. A weak electrolyte only partially dissociates — only some molecules or formula units produce ions. Weak acids and weak bases are the most common examples. In this chapter you are working primarily with strong electrolytes.
The concept of electrolytes appears throughout biology, medicine, and engineering. In the human body, dissolved ions — particularly Na⁺, K⁺, Ca²⁺, Mg²⁺, Cl⁻, and HCO₃⁻ — regulate fluid balance, nerve signaling, muscle contraction, and bone structure. These dissolved ions are electrolytes in the biochemical sense. Severe electrolyte imbalances — too little sodium, too little potassium — can be life-threatening.
In electrochemistry, electrolytes enable batteries and electrolysis cells to function by carrying charge between electrodes through ion movement. You will explore this in depth in Chapter 24. In water quality testing, the electrical conductivity of water gives a quick measurement of dissolved ion concentration.
The solid Earth is largely made of ionic compounds. Rocks and minerals — the materials that form mountains, ocean floors, and continental crust — are predominantly ionic or have ionic character.
| Mineral | Formula | Where Found / Uses |
|---|---|---|
| Halite | NaCl | Salt deposits, ocean evaporation; the source of most table salt |
| Calcite | CaCO₃ | Limestone, marble, chalk, coral reefs, mollusk shells, eggshells |
| Fluorite | CaF₂ | Cubic crystals; used in optical instruments and steel production |
| Gypsum | CaSO₄·2H₂O | Mined worldwide; used in plaster, drywall, and building materials |
| Corundum | Al₂O₃ | Pure form used as abrasive; ruby and sapphire are Al₂O₃ with trace ion impurities |
Pure aluminum oxide (Al₂O₃) is colorless. When trace amounts of chromium ions (Cr³⁺) replace a small fraction of aluminum ions in the lattice, the compound becomes ruby — deep red. When trace amounts of iron and titanium ions are present instead, it becomes sapphire — blue. The brilliant colors of precious gemstones are often due to small amounts of ionic impurities in otherwise colorless crystal lattices.
Many of the most important industrial chemicals are ionic compounds. Sodium hydroxide (NaOH) is one of the most widely produced chemicals in the world, used in manufacturing paper, textiles, soap, and detergents. Sodium carbonate (Na₂CO₃) is essential in making glass and ceramics. Ammonium nitrate (NH₄NO₃) is a common fertilizer providing nitrogen to crops. Calcium chloride (CaCl₂) is spread on icy roads in winter because it dissolves exothermically, releasing heat and lowering the freezing point of water. Potassium iodide (KI) is added to table salt in many countries to prevent iodine deficiency, which affects thyroid function.
Living things depend on ionic compounds and ions in solution for almost every biological function. Calcium phosphate — a combination of Ca²⁺ and PO₄³⁻ ions — is the major structural material in bones and teeth. The specific mineral hydroxyapatite gives the skeleton its hardness and rigidity.
Sodium and potassium ions in solution across cell membranes are the basis for nerve signals. When a nerve fires, Na⁺ ions rush into the cell and K⁺ ions rush out, creating an electrical signal that travels along the nerve fiber at high speed. Iron ions in hemoglobin carry oxygen in the blood — specifically Fe²⁺ in functioning hemoglobin. If iron is oxidized to Fe³⁺, hemoglobin cannot carry oxygen properly. Magnesium ions (Mg²⁺) are found at the center of chlorophyll, the molecule that captures sunlight in photosynthesis.
For each question, identify the ion that forms and its charge.
Write the correct ionic compound formula for each pair of ions. Show the ion charges and balancing process before writing the final formula. Then write the dissociation equation for each compound.
Answer these questions in complete sentences to demonstrate understanding, not just memorization.
For each property listed, explain it using ion charges, crystal lattice structure, and electrostatic attraction.
Each statement contains an error in reasoning. Identify the error and explain what is actually correct.
What is wrong with this statement? What are the actual particles in sodium chloride?
Where do the ions go when NaCl dissolves? Are they still present in the solution?
Why do ionic compounds actually require very high temperatures to melt? What must be overcome?
How is an ionic bond fundamentally different from a covalent bond? What happens to the electrons in each case?
What does the subscript 2 actually mean? What type of particles are present and in what ratio?
Give two examples of ionic compounds that do not dissolve readily in water and explain why solubility varies.
What does the formula unit actually describe? Is there an individual "NaCl particle" in a salt crystal?
Why does solid NaCl not conduct electricity? What must happen before an ionic compound can conduct?
Chapter 9 introduced ionic bonding — the transfer of electrons from metal atoms to nonmetal atoms and the resulting attraction between oppositely charged ions. Ionic bonding is one major way atoms achieve stable electron arrangements and form compounds. But what happens when two nonmetal atoms need electrons to become stable? Neither can simply give electrons to the other, because both atoms tend to attract electrons. Instead, they share electrons. This sharing of electrons between atoms is covalent bonding, and it produces molecules — discrete particles made of atoms held together by shared electron pairs. Chapter 10 introduces covalent bonding and molecular compounds. You will learn how electron pairs are shared between atoms, how Lewis structures represent bonds and lone pairs, how single, double, and triple bonds differ, and how the properties of molecular compounds compare to the ionic compounds you have just studied. The covalent world is the world of water, carbon dioxide, sugar, fats, proteins, plastics, and most of the compounds found in living things — making Chapter 10 one of the most important chapters in this textbook.