3.2 How Metals and Non-metals React: Ionic Compounds
Sodium is a soft metal that reacts violently with water; chlorine is a poisonous gas. Together they make the salt on your table. Learn why metals give away electrons, non-metals take them, and how the resulting ions lock together into crystals with high melting points that conduct only when molten or dissolved.
Drop a piece of sodium into water and it skates about, hissing and sometimes catching fire. Chlorine is a pale green gas so poisonous that it was once used as a weapon. Yet if you let the two react, the product is common salt, the white crystals you sprinkle on your food. How can two dangerous elements combine to make something we eat every day? The answer lies in what happens to their electrons — and the same idea explains magnesium oxide in antacids, calcium chloride in drying agents and thousands of other compounds that metals form with non-metals.
In this section you will see why atoms react at all, how a metal atom hands electrons to a non-metal atom to form ions, how the oppositely charged ions hold together as an ionic bond, and why ionic compounds have such characteristic properties. You will draw electron-shell diagrams for $\mathrm{NaCl}$, $\mathrm{MgCl_{2}}$ and $\mathrm{MgO}$ and learn to write the formula of an ionic compound from the charges of its ions. The ideas return in Class 11 Chemistry, Chapter 4 (Chemical Bonding). A 📘 box near the end, Periodic classification of elements, uses the same electron arrangements to build the periodic table; it is assessed in school but not in the year-end board exam.
Why Atoms React: the Stable Outer Shell
An atom has a tiny nucleus, with positively charged protons, surrounded by negatively charged electrons that move in shells. The shells are filled in order: the first shell (K) holds at most $2$ electrons, the second (L) at most $8$, the third (M) at most $8$ when it is the outermost one (the rule for the first twenty elements). The electrons in the outermost shell are the valence electrons; they decide how the atom reacts.
The noble gases — helium, neon, argon — are almost completely unreactive. Their outer shells are full: helium has $2$ electrons, neon $2,8$ and argon $2,8,8$. This full-shell arrangement is exceptionally stable. Every other atom reacts in a way that gives it the electron arrangement of the nearest noble gas.
| Element | Atomic number | Electrons in shells K, L, M, N | Valence electrons | Nearest noble gas |
|---|---|---|---|---|
| Lithium Li | 3 | 2, 1 | 1 | helium (2) |
| Oxygen O | 8 | 2, 6 | 6 | neon (2, 8) |
| Fluorine F | 9 | 2, 7 | 7 | neon (2, 8) |
| Neon Ne | 10 | 2, 8 | 8 | — |
| Sodium Na | 11 | 2, 8, 1 | 1 | neon (2, 8) |
| Magnesium Mg | 12 | 2, 8, 2 | 2 | neon (2, 8) |
| Aluminium Al | 13 | 2, 8, 3 | 3 | neon (2, 8) |
| Chlorine Cl | 17 | 2, 8, 7 | 7 | argon (2, 8, 8) |
| Argon Ar | 18 | 2, 8, 8 | 8 | — |
| Calcium Ca | 20 | 2, 8, 8, 2 | 2 | argon (2, 8, 8) |
Look at sodium, $2,8,1$. It is one electron more than neon. The easiest way to reach a full outer shell is to lose that single electron: the second shell, already full, becomes the outermost. Chlorine, $2,8,7$, is one electron short of argon, so the easiest way for it is to gain one electron. An atom of magnesium ($2,8,2$) loses two electrons; an atom of oxygen ($2,6$) gains two. Metals have few valence electrons (usually $1$, $2$ or $3$) and tend to lose them; non-metals have $4$ to $7$ and tend to gain.
Forming Ionic Compounds: NaCl, MgCl₂ and MgO
Now bring sodium and chlorine together. The sodium atom hands its single valence electron to the chlorine atom. Each becomes an ion with a full outer shell: $$\mathrm{Na \rightarrow Na^{+} + e^{-}}\qquad\qquad \mathrm{Cl + e^{-} \rightarrow Cl^{-}}$$ The electron moves from sodium to chlorine, which is shown below with the sodium electron as a dot and the chlorine electrons as crosses.
The two ions now have opposite charges, and opposite charges attract. The strong electrostatic force of attraction between a positive and a negative ion is the ionic bond.
Magnesium chloride. A magnesium atom ($2,8,2$) has two valence electrons to lose; a chlorine atom needs only one. So one magnesium atom gives one electron to each of two chlorine atoms: $$\mathrm{Mg \rightarrow Mg^{2+} + 2e^{-}}$$ $$\mathrm{Cl + e^{-} \rightarrow Cl^{-}}\quad(\text{twice})$$ The crystal contains one $\mathrm{Mg^{2+}}$ for every two $\mathrm{Cl^{-}}$: $+2-1-1=0$. The formula is $\mathrm{MgCl_{2}}$: $$\mathrm{Mg(s) + Cl_{2}(g) \rightarrow MgCl_{2}(s)}$$ Magnesium oxide. An oxygen atom ($2,6$) needs two electrons — exactly the two that magnesium offers. One $\mathrm{Mg^{2+}}$ and one $\mathrm{O^{2-}}$ balance, so the formula is $\mathrm{MgO}$: $$\mathrm{2Mg(s) + O_{2}(g) \rightarrow 2MgO(s)}$$ Sodium oxide needs the reverse ratio: oxygen takes two electrons but each sodium offers only one, so two Na atoms are needed per O atom: $\mathrm{2Na^{+}}$ and $\mathrm{O^{2-}}$ give $\mathrm{Na_{2}O}$.
Properties of Ionic Compounds
In a crystal of sodium chloride, every $\mathrm{Na^{+}}$ is surrounded by six $\mathrm{Cl^{-}}$ and every $\mathrm{Cl^{-}}$ by six $\mathrm{Na^{+}}$, in a regular three-dimensional pattern that repeats in every direction. There are no separate “$\mathrm{NaCl}$ molecules”; the whole crystal is one giant arrangement of ions held by strong attractions in all directions. From this picture most properties follow.
- Physical state: hard, brittle crystalline solids. The ions are locked in place, so the solid is rigid; but a blow slides one layer of ions over another so that like charges face each other, repel, and the crystal shatters.
- High melting and boiling points: a great deal of energy is needed to break the strong attractions between ions. Sodium chloride melts at $1074\ \mathrm{K}$ ($801\ ^{\circ}\mathrm{C}$) and boils at $1686\ \mathrm{K}$.
- Solubility: they are generally soluble in water (which pulls the ions apart) and insoluble in solvents such as kerosene and petrol.
- Conduction of electricity: they do not conduct in the solid state because the ions cannot move; they conduct when molten or dissolved in water, because the ions are then free to move and carry the current.
Compare three typical melting points: sodium chloride (ions $+1$ and $-1$) melts at $1074\ \mathrm{K}$, calcium chloride at $1045\ \mathrm{K}$, and magnesium oxide, whose ions carry charges $+2$ and $-2$ and attract each other much more strongly, at about $3125\ \mathrm{K}$ ($2852\ ^{\circ}\mathrm{C}$). Ionic compounds therefore stay solid at temperatures at which most metals and all molecular substances have melted.
Periodic classification of elements is part of the CBSE Class 10 syllabus and is taught and tested in school, but it is not examined in the year-end board exam. It belongs here because it rests on exactly the electron arrangements used in this section. The ideas return in Class 11 Chemistry, Ch. 3 — Classification of Elements and Periodicity in Properties.
Why classify? Well over a hundred elements are known. Learning each one separately would be hopeless, so chemists looked for a pattern that puts similar elements together. Three early attempts led to the table we use today.
1. Döbereiner's triads (1817–1829). Johann Döbereiner noticed groups of three elements with similar properties. Written in order of increasing atomic mass, the atomic mass of the middle element is roughly the average of the other two. Lithium, sodium and potassium ($6.9$, $23.0$ and $39.1\ \mathrm{u}$) form a triad: $\tfrac{6.9+39.1}{2}=23.0$. Calcium, strontium, barium and chlorine, bromine, iodine are two more. Limitation: he could find only three such triads; most of the elements then known did not fit.
2. Newlands' law of octaves (1866). John Newlands arranged the elements known to him in order of increasing atomic mass, starting with hydrogen, and found that every eighth element has properties similar to the first, like the eighth note of a musical scale (sa, re, ga, ma, pa, dha, ni, sa). Lithium is 2nd and sodium 9th; sodium is 9th and potassium 16th. Limitations: the pattern worked only up to calcium; he assumed only $56$ elements existed and none would be discovered; to make it fit he placed two elements in one slot (cobalt and nickel) in the same column as fluorine, chlorine and bromine, which are quite unlike them, while iron, which resembles cobalt and nickel, ended up far away. The noble gases, discovered later, could not be fitted at all.
3. Mendeleev's periodic table (1869). Dmitri Mendeleev took the $63$ elements then known and arranged them by increasing atomic mass, but grouped them above all by their chemical properties, especially the formulae of their oxides and hydrides. His table has eight vertical groups and horizontal periods, and rests on Mendeleev's periodic law: the properties of elements are a periodic function of their atomic masses.
- Achievements. Where no known element fitted, he left a gap and predicted the missing element. Eka-aluminium (predicted atomic mass $68$, oxide $\mathrm{E_{2}O_{3}}$) was found as gallium ($69.7\ \mathrm{u}$, $\mathrm{Ga_{2}O_{3}}$); eka-boron and eka-silicon became scandium and germanium. He put similar elements together even where it reversed the order of atomic masses, and the noble gases, discovered later, fitted in as a new group without disturbing the rest.
- Limitations. Hydrogen has no fixed place: like the alkali metals it forms compounds such as $\mathrm{HCl}$ (compare $\mathrm{NaCl}$), but like the halogens it exists as diatomic molecules, $\mathrm{H_{2}}$ (compare $\mathrm{Cl_{2}}$). Isotopes of an element have different atomic masses but the same properties, so ordering by mass gives them different places. Atomic masses do not rise in regular steps, so one cannot tell how many elements lie between two of them; and cobalt ($58.9\ \mathrm{u}$) has to come before nickel ($58.7\ \mathrm{u}$).
4. The Modern Periodic Table. In 1913 Henry Moseley showed that the atomic number $Z$ (the number of protons) is a more fundamental property than atomic mass. The Modern Periodic Law states: the properties of elements are a periodic function of their atomic numbers. Arranged by $Z$, the table has 18 groups (columns) and 7 periods (rows), and the problems of Mendeleev's table disappear: isotopes share one $Z$ and one place, and cobalt ($Z=27$) correctly comes before nickel ($Z=28$). Only hydrogen remains a special case. The position of an element follows directly from its electron arrangement:
- Elements in the same group have the same number of valence electrons, and therefore similar chemical properties. For the elements of this course, group $=$ valence electrons (groups 1, 2) or valence electrons $+\,10$ (groups 13 to 18); helium, with $2$, sits in group 18 with the other noble gases.
- Elements in the same period have the same number of occupied shells: period number $=$ number of shells. Period 1 has $2$ elements (the K shell holds $2$ electrons); periods 2 and 3 have $8$ each.
Trends in the Modern Periodic Table (main groups, noble gases left out of the size comparison):
| Property | Across a period (left → right) | Down a group |
|---|---|---|
| Valency | rises from $1$ to $4$, then falls to $0$: Na $1$, Mg $2$, Al $3$, Si $4$, P $3$, S $2$, Cl $1$, Ar $0$ | stays the same (same valence electrons) |
| Atomic size (radius) | decreases: Na $186$ → Cl $99\ \mathrm{pm}$ | increases: Li $152$, Na $186$, K $231\ \mathrm{pm}$ |
| Metallic character | decreases (metals on the left, non-metals on the right) | increases |
Valency is the number of electrons an atom loses, gains or shares to reach an octet: the number of valence electrons if it is $4$ or fewer, otherwise $8$ minus that number. Atomic size falls across a period because the nuclear charge grows while the new electrons enter the same shell, so they are pulled in more tightly; it rises down a group because each step adds a new shell. Metallic character is the tendency to lose electrons. It falls across a period (the larger nuclear charge holds the valence electrons more firmly) and rises down a group (the valence electrons are farther from the nucleus). The metalloids boron, silicon, germanium, arsenic, antimony and tellurium lie along a zig-zag line between metals and non-metals and have intermediate properties. Non-metallic character runs the opposite way, which is why the oxides change from basic on the left (e.g. $\mathrm{Na_{2}O}$) to acidic on the right (e.g. $\mathrm{SO_{2}}$), as in Section 3.1.
📘 Worked example 1 — Testing a triad. The atomic masses of calcium, strontium and barium are $40.1$, $87.6$ and $137.3\ \mathrm{u}$. Do they form a Döbereiner triad? Where do the three elements stand in the Modern Periodic Table?
Solution. The average of the first and third masses is $\tfrac{40.1+137.3}{2}=88.7\ \mathrm{u}$, close to strontium's $87.6\ \mathrm{u}$, and the three elements are chemically alike (reactive metals forming $\mathrm{M^{2+}}$ ions), so yes, they form a triad. All three have $2$ valence electrons: they lie together in group 2, in periods 4, 5 and 6, with atomic size and metallic character increasing from calcium to barium.
📘 Worked example 2 — Placing elements. Elements A, B and C have atomic numbers $11$, $14$ and $17$. (a) Write their electron arrangements and find the group and period of each. (b) Give the valency of each. (c) Which has the largest atoms, and which is the most metallic? (d) Write the formula of the compound of A and C.
Solution. (a) A: $2,8,1$, group 1; B: $2,8,4$, group $4+10=14$; C: $2,8,7$, group $7+10=17$. All have three shells, so all are in period 3. (b) A: $1$; B: $4$; C: $8-7=1$. (c) Size decreases across a period, so A is the largest; A, on the far left, is also the most metallic (it is sodium; B is silicon, a metalloid; C is chlorine, a non-metal). (d) A gives one electron to C: $\mathrm{A^{+}C^{-}}$, formula AC — sodium chloride.
1. “Salt conducts electricity.” Solid salt does not. It conducts only when molten or dissolved, because only then are the ions free to move.
2. Writing $\mathrm{NaCl}$ as a molecule. An ionic compound is a lattice of ions; $\mathrm{NaCl}$ is a formula unit giving the ratio $1:1$.
3. Wrong charges in the formula. $\mathrm{Al_{2}O_{3}}$ is $2\times(+3)+3\times(-2)=0$, not $\mathrm{AlO}$. Balance the charges, not the atoms of a diagram.
4. Mixing up cation and anion. Cations are positive because they have lost electrons; anions are negative because they have gained them. Metals form cations, non-metals form anions.
5. Changing the nucleus. Gaining or losing electrons never changes the number of protons.
Exercises
- 1 Write the electron arrangement (K, L, M) of Li, O, F, Na, Mg, Al and Cl.
- 2 Why are the noble gases almost completely unreactive?
- 3 Define an ion, a cation and an anion, with one example of each.
- 4 An element has atomic number $13$. Write its electron arrangement, the ion it forms and the noble gas whose arrangement the ion has.
- 5 Write the symbol (with charge) of the ion formed by (a) oxygen, (b) fluorine, (c) magnesium, (d) sodium, (e) chlorine.
- 6 Why does sodium lose an electron rather than gain seven?
- 7 Give the number of protons and electrons in (a) $\mathrm{Cl^{-}}$ ($Z=17$), (b) $\mathrm{Mg^{2+}}$ ($Z=12$), (c) $\mathrm{O^{2-}}$ ($Z=8$).
- 8 Do the electrons or the protons change when a sodium atom becomes a sodium ion? How does the size of the charge arise?
- 1 Describe, with equations for each atom, how sodium chloride forms from sodium and chlorine.
- 2 Why do two chlorine atoms combine with one magnesium atom in $\mathrm{MgCl_{2}}$?
- 3 Show the formation of magnesium oxide from a magnesium atom and an oxygen atom, giving the charges on the ions.
- 4 Write the formula of (a) sodium oxide, (b) calcium chloride, (c) aluminium oxide, (d) aluminium chloride, (e) magnesium oxide.
- 5 Write a balanced equation with state symbols for the formation of calcium chloride from calcium and chlorine.
- 6 What mass of sodium chloride is formed when $2.3\ \mathrm{g}$ of sodium reacts completely with chlorine? ($\mathrm{Na}=23\ \mathrm{u}$, $\mathrm{Cl}=35.5\ \mathrm{u}$)
- 7 Find the formula-unit mass of $\mathrm{MgCl_{2}}$ and the percentage of magnesium in it. ($\mathrm{Mg}=24\ \mathrm{u}$, $\mathrm{Cl}=35.5\ \mathrm{u}$)
- 8 Calcium (atomic number $20$) reacts with oxygen. Give its electron arrangement, the ions formed, and the formula and equation of the product.
- 9 Why is the formula of aluminium oxide $\mathrm{Al_{2}O_{3}}$ and not $\mathrm{AlO}$?
- 1 Why are ionic solids hard but also brittle?
- 2 Why do ionic compounds have high melting points?
- 3 Why does sodium chloride conduct electricity when molten or dissolved in water but not when solid?
- 4 In which of these liquids does sodium chloride dissolve: water, kerosene, petrol?
- 5 A bulb is connected in series with electrodes. Which of the following make it glow: solid KCl, molten KCl, KCl solution, sugar solution?
- 6 A compound has a high melting point, is brittle, and its solution conducts electricity. A second compound is a soft solid that melts at $330\ \mathrm{K}$, does not dissolve in water and does not conduct. Which is likely ionic?
- 7 The melting points of sodium chloride, calcium chloride and magnesium oxide are $1074\ \mathrm{K}$, $1045\ \mathrm{K}$ and about $3125\ \mathrm{K}$. Convert them to degrees Celsius and suggest why magnesium oxide is so high.
- 8 Applied. Pure distilled water hardly lights a bulb, but when a little common salt is dissolved in it the bulb glows brightly. Explain. Why are electrical appliances never touched with wet hands?
- 1 Element X has electron arrangement $2,8,3$ and element Y has $2,6$. Name the type of bond between them and write the formula of the compound.
- 2 Potassium has atomic number $19$ ($2,8,8,1$). Write the formulae of potassium chloride and potassium oxide and balanced equations for their formation from the elements.
- 3 Compare $\mathrm{Na^{+}}$ and a neon atom: what do they have in common and how do they differ?
- 4 Why can two metals not form an ionic compound with each other?
- 5 Challenging. Molten sodium chloride is electrolysed between two electrodes. Towards which electrode does each ion move, and what is formed at each? Write the overall equation.
- 1 Lithium, sodium and potassium have atomic masses $6.9$, $23.0$ and $39.1\ \mathrm{u}$. Show that they form a Döbereiner triad.
- 2 State Newlands' law of octaves and give two of its limitations.
- 3 Give two achievements and two limitations of Mendeleev's periodic table.
- 4 State the Modern Periodic Law. How does it settle the problem of where to place isotopes?
- 5 An element has atomic number $16$. Find its electron arrangement, group, period and valency, and say whether it is a metal or a non-metal.
- 6 Arrange $\mathrm{Mg}$, $\mathrm{Na}$, $\mathrm{Cl}$ and $\mathrm{P}$ in order of decreasing atomic size, and say which is the most metallic.
E1. Board · 1 mark Which of the following is not a property of ionic compounds? (A) high melting point (B) conduct electricity in the molten state (C) soluble in water (D) conduct electricity in the solid state
E2. Board · A–R Assertion (A): Sodium chloride conducts electricity in the molten state but not in the solid state. Reason (R): In the molten state the ions are free to move. Choose: (A) Both A and R are true and R is the correct explanation of A. (B) Both A and R are true but R is not the correct explanation of A. (C) A is true but R is false. (D) A is false but R is true.
E3. Board · 3 marks The atomic numbers of elements A and B are $12$ and $17$. (a) Write their electron arrangements. (b) Write the formula of the compound formed and name the type of bond. (c) Show the transfer of electrons.
E4. Board · Case-based Three solids are tested. A: melting point $1074\ \mathrm{K}$, dissolves in water, solution conducts, solid does not. B: melting point $340\ \mathrm{K}$, insoluble in water, soluble in petrol, does not conduct in any state. C: melting point $1356\ \mathrm{K}$ (copper), insoluble, conducts as a solid. (i) Which is ionic, and which is a metal? (2 marks) (ii) Why does A not conduct when solid? (1 mark) (iii) If A is sodium chloride, write the equation for the formation of its ions in water. (1 mark)
E5. JEE Main / NEET Foundation. Of the ions $\mathrm{Na^{+}}$, $\mathrm{Mg^{2+}}$, $\mathrm{F^{-}}$, $\mathrm{O^{2-}}$ and $\mathrm{Cl^{-}}$, how many have the electron arrangement $2,8$? (A) 3 (B) 4 (C) 5 (D) 2
E6. Board · 2 marks Why do ionic compounds have high melting points and why are they soluble in water but not in kerosene?
Research question. Which household solutions carry an electric current, and how does the concentration of dissolved salt change the brightness of an LED in a home-made conductivity tester?
Why it matters. Conductivity meters are used all over the world to test whether drinking water, river water and soil water are too salty. Hospitals rely on carefully balanced ionic solutions in drips, sports drinks replace salts lost in sweat, and the electrical signals in your nerves are carried by moving $\mathrm{Na^{+}}$ and $\mathrm{K^{+}}$ ions. A simple tester shows that the more ions in water, the better it conducts.
What to do.
- Build a tester: two graphite pencil leads (or two clean iron nails) held apart in a cork, connected in series with a $4.5\ \mathrm{V}$ battery and an LED with a suitable resistor (ask a teacher).
- Test: distilled or boiled and cooled water, tap water, sugar solution, salt solution, baking-soda solution, vinegar, lemon juice and cooking oil. Record whether the LED glows and how brightly.
- Make salt solutions with $1$, $2$, $4$ and $8$ teaspoons of salt in a glass of water and rank the brightness.
- Dissolve and test a pinch of washing soda and a pinch of sugar for comparison.
Tools: Battery, LED, resistor, wires, pencil leads, plastic cups, salt, sugar, and household liquids.
Hand in: A table of substances with brightness (none / dim / bright), a graph of brightness against salt added, and a paragraph explaining the results in terms of ions.
Going further. Does brightness keep rising without limit? Does the LED glow for solid salt pressed between the electrodes?
🧪 Self-check quiz
Q1. The electron arrangement of a magnesium ion, $\mathrm{Mg^{2+}}$, is
Q2. Which of the following is the correct formula of aluminium oxide?
Q3. Solid sodium chloride does not conduct electricity because
Q4. In the formation of $\mathrm{MgO}$ from Mg and O, how many electrons does each magnesium atom give away?
Q5. Which statement about ionic compounds is correct?