A Poisonous Gas and an Explosive Metal Combine Into the Salt on Your Table

Sodium is a soft metal that reacts violently with water. Chlorine is a greenish-yellow gas poisonous enough to damage the lungs. Put them together and the product is sodium chloride — the salt sprinkled on food every day.

Something drastic must have happened to both atoms, and it has. Sodium has given one electron to chlorine.

- Sodium, $2,8,1$, loses its single outer electron and becomes $2,8$ — the complete octet of neon
- Chlorine, $2,8,7$, gains that electron and becomes $2,8,8$ — the complete octet of argon

Both now have the stable arrangement they lacked, so neither has any further tendency to react. That is why the compound is so safe while the elements are so dangerous.

But the electron transfer leaves the two atoms charged. Sodium, having lost a negative electron, carries a charge of $+1$; chlorine, having gained one, carries $-1$. Opposite charges attract strongly, and that electrostatic attraction between the two ions is the electrovalent bond, also called the ionic bond.

From that picture the whole of this part follows.

- Which elements form such bonds — metals that lose electrons easily with non-metals that gain them readily
- How to show the transfer for a range of compounds, including ones where the charges are not equal
- Why ionic compounds behave as they do — hard, high-melting, soluble in water, and conducting electricity only when their ions are free to move
- Why that last property makes electrolysis possible, which connects this chapter directly to the chapter on electrolysis

One idea is worth keeping in view throughout. An ionic compound is not a collection of separate molecules. It is a vast regular arrangement of positive and negative ions, each held by all its neighbours — and almost every property of the compound comes from that arrangement.

This page covers the first part of the ICSE Class 10 Chemistry chapter on chemical bonding: the electrovalent bond, electron dot structures, the properties of electrovalent compounds, and their dissociation and electrolysis.

A metal atom transfers one or more electrons to a non-metal atom, both reach a stable octet or duplet, and the oppositely charged ions formed attract each other.

The definition. An electrovalent bond is the chemical bond formed between two atoms by the complete transfer of one or more electrons from one atom to the other, as a result of which both attain a stable electronic configuration. The electrostatic force of attraction between the resulting ions holds them together.

Why atoms transfer electrons at all. Atoms other than the noble gases have incomplete outer shells, and they react to reach the configuration of the nearest noble gas.

- An atom with one, two or three outer electrons finds it easier to lose them than to gain five, six or seven. It becomes a positive ion, a cation
- An atom with five, six or seven outer electrons finds it easier to gain three, two or one. It becomes a negative ion, an anion

Electrovalency. The number of electrons an atom loses or gains during the formation of an electrovalent bond is its electrovalency.

- Sodium loses one electron, so its electrovalency is $+1$
- Magnesium loses two, so its electrovalency is $+2$
- Chlorine gains one, so its electrovalency is $-1$
- Oxygen gains two, so its electrovalency is $-2$

The conditions that favour an electrovalent bond, which link this chapter to the previous one on periodic properties:

- The metal atom should have a low ionisation potential, so that it loses electrons easily. Alkali and alkaline earth metals qualify best
- The non-metal atom should have a high electron affinity, so that it gains electrons readily. Halogens and oxygen qualify best
- The difference in electronegativity between the two atoms should be large

So an electrovalent bond forms most readily between elements from the left of the table and elements from the right — the two ends that the periodic trends made most different.

Worked example — the steps for sodium chloride, written as equations.

Sodium loses an electron:

$$\text{Na}\to {\text{Na}}^{+}+e^{-}$$
$$2,8,1\to 2,8$$

Chlorine gains it:

$$\text{Cl}+e^{-}\to {\text{Cl}}^{-}$$
$$2,8,7\to 2,8,8$$

The ions attract:

$${\text{Na}}^{+}+{\text{Cl}}^{-}\to \text{NaCl}$$

Check the electrons and the charges. One electron was lost and one gained, so the electrons balance; $+1$ and $-1$ make zero, so the compound is neutral. Every electrovalent compound must be electrically neutral, and that requirement is what fixes the ratio of the ions in the formula.

Why the formula shows a ratio and not a molecule. In solid sodium chloride each sodium ion is surrounded by six chloride ions and each chloride ion by six sodium ions, in a repeating three-dimensional pattern called a crystal lattice. No single sodium ion belongs to any particular chloride ion, so there is no discrete molecule of sodium chloride. **The formula $\text{NaCl}$ simply records that the ions are present in the ratio one to one.

One boundary case worth stating. A metal atom and a non-metal atom do not always form an electrovalent bond. When the difference in electronegativity is small, electrons are shared instead of transferred, giving a covalent bond — the subject of the next part of the chapter. The electrovalent bond is the extreme case**, found when one atom holds its electrons weakly and the other attracts them strongly.

For each compound, show each metal atom losing its outer electrons, each non-metal atom gaining enough to complete its octet, and use as many atoms of each as are needed to balance the electrons transferred.

The convention in a dot structure. Outer electrons are shown as dots round the symbol of one atom and as crosses round the other, so that the transferred electrons can be followed. After transfer, the ions are written in square brackets with their charges, and the non-metal ion is shown with a complete octet of dots and crosses.

**1. Sodium chloride, $\text{NaCl}$.

- Na** $2,8,1$ gives its one outer electron, shown as a cross, to Cl $2,8,7$
- Result: $[\text{Na}{]}^{+}$ and $[\text{Cl}{]}^{-}$, the chloride ion carrying seven dots and one cross — eight in all
- **Ratio $1:1$, since one electron lost matches one gained

2. Magnesium chloride, ${\text{MgCl}}_2$.**

$$\text{Mg}\to {\text{Mg}}^{2+}+2e^{-},2\text{Cl}+2e^{-}\to 2{\text{Cl}}^{-}$$

- Mg $2,8,2$ loses two electrons to become $2,8$
- Each Cl $2,8,7$ can accept only one, so two chlorine atoms are needed
- Result: $[\text{Mg}{]}^{2+}$ and two $[\text{Cl}{]}^{-}$ ions, **ratio $1:2$

3. Calcium oxide, $\text{CaO}$.**

$$\text{Ca}\to {\text{Ca}}^{2+}+2e^{-},\text{O}+2e^{-}\to {\text{O}}^{2-}$$

- Ca $2,8,8,2$ loses two electrons to become $2,8,8$
- O $2,6$ gains both to become $2,8$
- Result: $[\text{Ca}{]}^{2+}$ and $[\text{O}{]}^{2-}$, **ratio $1:1$ — two electrons lost and two gained

4. Potassium oxide, ${\text{K}}_2\text{O}$.**

$$2\text{K}\to 2{\text{K}}^{+}+2e^{-},\text{O}+2e^{-}\to {\text{O}}^{2-}$$

- Each K $2,8,8,1$ can give only one electron, becoming $2,8,8$
- O needs two, so two potassium atoms are needed
- Result: two $[\text{K}{]}^{+}$ and one $[\text{O}{]}^{2-}$, **ratio $2:1$

5. Sodium oxide, ${\text{Na}}_2\text{O}$.**

$$2\text{Na}\to 2{\text{Na}}^{+}+2e^{-},\text{O}+2e^{-}\to {\text{O}}^{2-}$$

- The same pattern as potassium oxide — two sodium atoms, each giving one electron, to one oxygen atom needing two
- Result: two $[\text{Na}{]}^{+}$ and one $[\text{O}{]}^{2-}$, **ratio $2:1$

6. Sodium nitride, ${\text{Na}}_3\text{N}$.**

$$3\text{Na}\to 3{\text{Na}}^{+}+3e^{-},\text{N}+3e^{-}\to {\text{N}}^{3-}$$

- N $2,5$ needs three electrons to reach $2,8$
- Each Na supplies only one, so three sodium atoms are needed
- Result: three $[\text{Na}{]}^{+}$ and one $[\text{N}{]}^{3-}$, **ratio $3:1$

The check to apply to every one of them, which catches every error:

- Electrons lost must equal electrons gained
- Total positive charge must equal total negative charge
- Every ion must end with a complete octet — or a duplet if it has only one shell

Worked check — magnesium chloride.** Magnesium loses $2$ electrons; two chlorine atoms gain $1$ each, total $2$. Balanced. Charges: $+2$ from magnesium, $-1\times 2=-2$ from chlorine. Neutral. Configurations: magnesium ion $2,8$, each chloride ion $2,8,8$. All complete.

Worked check — sodium nitride. Three sodium atoms lose $3$ electrons; one nitrogen gains $3$. Charges $+3$ and $-3$. Neutral, and the nitride ion $2,8$ has a complete octet.

The pattern behind all six ratios. The ratio of atoms is the ratio that makes the electrons balance, which is the same as exchanging the electrovalencies — exactly the criss-cross rule for writing formulae. **Magnesium $2$ with chlorine $1$ gives ${\text{MgCl}}_2$; sodium $1$ with nitrogen $3$ gives ${\text{Na}}_3\text{N}$.** The dot structure is the reason the rule works.

Because they are built of oppositely charged ions held firmly in a lattice by strong electrostatic forces — strong enough to make the solid hard and high-melting, and firm enough to stop the ions moving until the lattice is broken.

1. State of existence — crystalline solids. Electrovalent compounds are solids at room temperature, with their ions arranged in a regular crystal lattice. The strong attraction in every direction holds the ions in fixed positions, which is why no ionic compound is a gas or a liquid under ordinary conditions.

2. Hard but brittle. The strong forces make the crystals hard. But they are also brittle, and the reason is worth stating.

- If a layer of ions is pushed sideways by one position, ions of the same charge come to lie opposite each other
- Like charges repel strongly, so the crystal splits along that plane
- So a crystal of common salt shatters when struck rather than bending

3. High melting and boiling points. Melting requires the ions to be pulled out of their fixed positions, which means overcoming very strong electrostatic attractions in every direction. A large amount of energy is needed, so electrovalent compounds melt and boil at high temperatures — sodium chloride stays solid far above the temperatures at which wax or sugar melt.

4. Soluble in water, insoluble in organic solvents.

- Water molecules are polar, carrying a slight negative charge on the oxygen and slight positive charges on the hydrogens
- They cluster round the ions — the oxygen ends toward positive ions and the hydrogen ends toward negative ones — and pull them out of the lattice
- So ionic compounds dissolve in water
- Organic solvents such as kerosene and petrol are non-polar, cannot attract the ions, and do not dissolve them

5. Conductivity of electricity — the property examined most often.

- In the solid state, electrovalent compounds do not conduct. The ions are present but fixed in the lattice, and there are no free electrons, so nothing can move to carry a current
- In the molten state, they conduct. Melting breaks down the lattice and the ions become free to move
- In aqueous solution, they conduct. Dissolving separates the ions and they are free to move through the water

So an ionic compound conducts only when its ions are mobile. The charge carriers are the ions themselves, not electrons — which is the essential difference from a metal.

6. Conductivity of heat. In the solid state electrovalent compounds are poor conductors of heat, since they have no free electrons to carry energy through the solid.

7. Non-volatile. Because of the strong forces, they have very low vapour pressures and do not evaporate at ordinary temperatures.

8. Fast reactions in solution. Reactions between ionic compounds in solution are very fast, because the ions are already present and only need to meet. Adding silver nitrate solution to sodium chloride solution gives a white precipitate of silver chloride instantly, since silver ions and chloride ions combine the moment they meet:

$${\text{AgNO}}_3+\text{NaCl}\to \text{AgCl}+{\text{NaNO}}_3$$

Worked comparison — the same solid in three states. Solid sodium chloride, molten sodium chloride and sodium chloride solution are each placed in a circuit with a bulb.

- Solid — the bulb does not glow, since the ions are fixed
- Molten — the bulb glows, since the ions are free
- Solution — the bulb glows, since the ions are free

Three states of one compound, and only the mobility of the ions changed.

One boundary case about brittleness and conductivity. A metal is also held together by strong forces, yet it bends rather than shattering and conducts even as a solid. The difference is the free electrons. In a metal, electrons are free to move and a sliding layer meets the same surroundings; in an ionic solid there are no free electrons and a sliding layer brings like charges together. One structural difference explains two opposite properties.

Melting or dissolving frees the ions that were already present in the solid, and free ions can move to the electrodes and be discharged — which is electrolysis.

Dissociation. The separation of the ions of an electrovalent compound when it is melted or dissolved in water is called dissociation.

$$\text{NaCl}\to {\text{Na}}^{+}+{\text{Cl}}^{-}$$

The ions are not created at that moment — they existed in the solid lattice all along. Melting or dissolving only releases them from their fixed positions. That is the distinction between dissociation, which applies to ionic compounds, and ionisation, in which a covalent compound such as hydrogen chloride actually forms ions for the first time when it dissolves in water.

An electrolyte. A compound that conducts electricity in the molten state or in aqueous solution and is decomposed by the current is an electrolyte. All electrovalent compounds are electrolytes, since all of them dissociate into mobile ions.

What happens when a current is passed through molten sodium chloride. Two electrodes connected to a battery are placed in the melt.

- The cathode is the negative electrode, connected to the negative terminal; the anode is the positive electrode
- Positive sodium ions move toward the cathode, attracted by its negative charge
- Negative chloride ions move toward the anode

At the cathode, each sodium ion gains an electron and becomes a sodium atom:

$${\text{Na}}^{+}+e^{-}\to \text{Na}$$

At the anode, chloride ions each lose an electron, and the chlorine atoms pair up as chlorine gas:

$$2{\text{Cl}}^{-}\to {\text{Cl}}_2+2e^{-}$$

Check the charges on the anode equation. Left side: $2\times (-1)=-2$. Right side: $0+2\times (-1)=-2$. Balanced.

So molten sodium chloride is decomposed into sodium metal and chlorine gas — the reverse of the reaction that formed it, driven by electrical energy.

Why the electrons balance overall. For every chloride ion that gives an electron to the anode, a sodium ion takes one from the cathode. The electrons given up at the anode travel round the external circuit to the cathode, so the current is carried by electrons in the wires and by moving ions in the melt.

Worked example — predicting the products. Name the products at each electrode when molten calcium chloride, ${\text{CaCl}}_2$, is electrolysed.

$${\text{CaCl}}_2\to {\text{Ca}}^{2+}+2{\text{Cl}}^{-}$$

- At the cathode, calcium ions each gain two electrons: ${\text{Ca}}^{2+}+2e^{-}\to \text{Ca}$. Calcium metal is formed
- At the anode, chloride ions lose electrons: $2{\text{Cl}}^{-}\to {\text{Cl}}_2+2e^{-}$. Chlorine gas is given off

Check: two electrons are taken by each calcium ion and two are released for every chlorine molecule, so the electrons balance.

Why reactive metals are extracted this way. Sodium, potassium, calcium and magnesium hold their electrons so weakly in the atom — and so reluctantly take them back as ions — that no ordinary chemical reducing agent can turn their ions into metal. Electrolysis of their molten compounds supplies the electrons directly at the cathode, which is why these metals are obtained by electrolysis.

A boundary case about solutions. When an ionic compound is electrolysed in aqueous solution rather than molten, the water also takes part, and the products at the electrodes can be different. Electrolysing sodium chloride solution does not give sodium at the cathode, because hydrogen from the water is discharged in preference. The electrolysis chapter explains which ions are discharged and why, but the starting point — free, mobile ions from dissociation — is exactly the one established here.