One Atom Hands Over an Electron and the Other Just Shares
Build electrovalent bonds by transferring electrons in sodium chloride, magnesium chloride and calcium oxide, draw covalent bonds for nine molecules, classify single, double and triple bonds, and compare the two kinds of compound.
Why do some atoms hand electrons over and others share them?
The previous part of this chapter ended with the octet rule: an atom combines in order to end up with eight electrons in its outermost shell, or two in the case of the first shell.
That tells you the destination. It does not tell you how an atom gets there — and there are only two routes.
Either an electron is handed over completely, or it is shared.
Sodium has one valence electron it would rather be rid of. Chlorine is one short of eight. So sodium hands its electron to chlorine, both reach an octet, and the two atoms are left as oppositely charged ions that attract one another. That is an electrovalent or ionic bond.
Now take two chlorine atoms. Each is one electron short, and neither has a spare to give. Handing one over would leave one of them worse off than before. So instead each contributes one electron to a shared pair that counts towards both octets at once. That is a covalent bond.
Transfer works when one atom wants to lose and the other wants to gain; sharing is what happens when both want to gain. Metals lose easily and non-metals gain, so a metal with a non-metal gives an ionic bond and two non-metals give a covalent one.
This page covers electrovalent bonding with worked orbit structures, covalent bonding across nine molecules, single, double and triple bonds, a full comparison of the two kinds of compound, and what happens when a covalent bond breaks.
That tells you the destination. It does not tell you how an atom gets there — and there are only two routes.
Either an electron is handed over completely, or it is shared.
Sodium has one valence electron it would rather be rid of. Chlorine is one short of eight. So sodium hands its electron to chlorine, both reach an octet, and the two atoms are left as oppositely charged ions that attract one another. That is an electrovalent or ionic bond.
Now take two chlorine atoms. Each is one electron short, and neither has a spare to give. Handing one over would leave one of them worse off than before. So instead each contributes one electron to a shared pair that counts towards both octets at once. That is a covalent bond.
Transfer works when one atom wants to lose and the other wants to gain; sharing is what happens when both want to gain. Metals lose easily and non-metals gain, so a metal with a non-metal gives an ionic bond and two non-metals give a covalent one.
This page covers electrovalent bonding with worked orbit structures, covalent bonding across nine molecules, single, double and triple bonds, a full comparison of the two kinds of compound, and what happens when a covalent bond breaks.
How does an electrovalent bond form in NaCl, MgCl2 and CaO?
An electrovalent bond is formed by the complete transfer of one or more electrons from a metal atom to a non-metal atom, so that both acquire a stable octet, and the oppositely charged ions produced are then held together by electrostatic attraction.
Three things happen in order: the metal loses, the non-metal gains, and the resulting ions attract.
Sodium chloride. Sodium is 2, 8, 1 and chlorine is 2, 8, 7.
- Sodium loses its single valence electron: (2, 8, 1) (2, 8), the neon arrangement
- Chlorine gains that electron: (2, 8, 7) (2, 8, 8), the argon arrangement
- and attract, giving
One electron moves, and both atoms are satisfied by the same transfer. That is why the formula is a simple one-to-one .
Magnesium chloride. Magnesium is 2, 8, 2 and needs to lose two electrons; each chlorine can take only one.
- (2, 8, 2) (2, 8), losing two electrons
- Two chlorine atoms are needed, each gaining one: (2, 8, 7) (2, 8, 8)
- One and two give
The subscript 2 is bookkeeping on the electrons. Magnesium has two to give away and each chlorine takes exactly one, so two chlorines are required — the formula is decided by the electron count, not chosen.
Calcium oxide. Calcium is 2, 8, 8, 2 and oxygen is 2, 6.
- (2, 8, 8, 2) (2, 8, 8), losing two electrons
- (2, 6) (2, 8), gaining two electrons
- and give
Here the two electrons calcium loses are exactly the two oxygen wants, so one atom of each is enough and the formula is with no subscripts. **Compare it with — the difference in formula comes entirely from how many electrons the non-metal can accept.
What the orbit structure must show. When a question asks for the structure of an electrovalent compound, draw the shells of both atoms, mark the transferred electron with an arrow from donor to acceptor, and write the charge on each ion. The ions are drawn separately, not joined by a line — there is no shared pair to draw.
An ionic compound has no molecules. Solid sodium chloride is a giant lattice** of alternating and ions extending in all three directions — there is no single particle you can point to as a molecule of . So the formula gives the ratio of ions rather than the contents of a molecule, which is why an ionic formula is properly called a formula unit, and why ionic compounds behave so differently from covalent ones in the comparison below.
Three things happen in order: the metal loses, the non-metal gains, and the resulting ions attract.
Sodium chloride. Sodium is 2, 8, 1 and chlorine is 2, 8, 7.
- Sodium loses its single valence electron: (2, 8, 1) (2, 8), the neon arrangement
- Chlorine gains that electron: (2, 8, 7) (2, 8, 8), the argon arrangement
- and attract, giving
One electron moves, and both atoms are satisfied by the same transfer. That is why the formula is a simple one-to-one .
Magnesium chloride. Magnesium is 2, 8, 2 and needs to lose two electrons; each chlorine can take only one.
- (2, 8, 2) (2, 8), losing two electrons
- Two chlorine atoms are needed, each gaining one: (2, 8, 7) (2, 8, 8)
- One and two give
The subscript 2 is bookkeeping on the electrons. Magnesium has two to give away and each chlorine takes exactly one, so two chlorines are required — the formula is decided by the electron count, not chosen.
Calcium oxide. Calcium is 2, 8, 8, 2 and oxygen is 2, 6.
- (2, 8, 8, 2) (2, 8, 8), losing two electrons
- (2, 6) (2, 8), gaining two electrons
- and give
Here the two electrons calcium loses are exactly the two oxygen wants, so one atom of each is enough and the formula is with no subscripts. **Compare it with — the difference in formula comes entirely from how many electrons the non-metal can accept.
What the orbit structure must show. When a question asks for the structure of an electrovalent compound, draw the shells of both atoms, mark the transferred electron with an arrow from donor to acceptor, and write the charge on each ion. The ions are drawn separately, not joined by a line — there is no shared pair to draw.
An ionic compound has no molecules. Solid sodium chloride is a giant lattice** of alternating and ions extending in all three directions — there is no single particle you can point to as a molecule of . So the formula gives the ratio of ions rather than the contents of a molecule, which is why an ionic formula is properly called a formula unit, and why ionic compounds behave so differently from covalent ones in the comparison below.
How does sharing work in hydrogen, oxygen, water and methane?
A covalent bond is formed by the mutual sharing of one or more pairs of electrons between two atoms, each atom contributing an equal number of electrons to the shared pair or pairs, so that both attain a stable octet or duplet.
A shared pair counts towards the octet of both atoms at once — which is precisely why sharing solves a problem transfer cannot.
**Hydrogen, . Each hydrogen has one electron and needs two — a duplet, since the first shell holds only two. Each contributes one electron to a single shared pair, and both reach the helium arrangement. One shared pair, a single bond**, written .
**Chlorine, . Each chlorine is 2, 8, 7 and needs one more. Each contributes one electron to one shared pair, giving both an octet. One shared pair, a single bond**, . Each chlorine also keeps three lone pairs that take no part in the bond.
**Oxygen, . Each oxygen is 2, 6 and needs two more. One shared pair is not enough, so each contributes two electrons and two pairs are shared. A double bond**, .
**Nitrogen, . Each nitrogen is 2, 5 and needs three. Each contributes three electrons and three pairs are shared. A triple bond**, .
So the number of shared pairs is just how many electrons each atom was short of. Oxygen short by two gives a double bond; nitrogen short by three gives a triple bond. Count the shortfall and you have the bond — there is nothing to memorise separately.
**Hydrogen chloride, . Hydrogen needs one and chlorine needs one. One shared pair — a single bond — leaves hydrogen with a duplet and chlorine with an octet.
Water, . Oxygen (2, 6) needs two electrons, and each hydrogen can supply one. So oxygen forms two single bonds, one to each hydrogen. Oxygen ends with an octet — two bond pairs and two lone pairs — and each hydrogen with a duplet.
Ammonia, . Nitrogen (2, 5) needs three, so it forms three single bonds to three hydrogens. Nitrogen ends with an octet — three bond pairs and one lone pair.
Methane, . Carbon (2, 4) needs four, so it forms four single bonds to four hydrogens. Carbon ends with an octet and no lone pairs.
Carbon tetrachloride, . Carbon forms four single bonds, one to each chlorine. Carbon reaches an octet, and each chlorine reaches an octet with three lone pairs of its own.
Water, ammonia and methane are the same idea with different shortfalls. Oxygen is short two and takes two hydrogens; nitrogen short three takes three; carbon short four takes four. So the formula , , is read straight off the valence-electron count, and drawing the structure is a matter of arranging bond pairs and then filling in whatever lone pairs are left over.
Counting the electrons in a structure. For water, oxygen brings 6 valence electrons and the two hydrogens bring 1 each, giving 8 in the picture: 2 bond pairs (4 electrons) and 2 lone pairs (4 electrons). If your drawing has more or fewer electrons than the atoms brought, it is wrong** — and this check catches the usual error of adding a lone pair that has nowhere to come from.
A shared pair counts towards the octet of both atoms at once — which is precisely why sharing solves a problem transfer cannot.
**Hydrogen, . Each hydrogen has one electron and needs two — a duplet, since the first shell holds only two. Each contributes one electron to a single shared pair, and both reach the helium arrangement. One shared pair, a single bond**, written .
**Chlorine, . Each chlorine is 2, 8, 7 and needs one more. Each contributes one electron to one shared pair, giving both an octet. One shared pair, a single bond**, . Each chlorine also keeps three lone pairs that take no part in the bond.
**Oxygen, . Each oxygen is 2, 6 and needs two more. One shared pair is not enough, so each contributes two electrons and two pairs are shared. A double bond**, .
**Nitrogen, . Each nitrogen is 2, 5 and needs three. Each contributes three electrons and three pairs are shared. A triple bond**, .
So the number of shared pairs is just how many electrons each atom was short of. Oxygen short by two gives a double bond; nitrogen short by three gives a triple bond. Count the shortfall and you have the bond — there is nothing to memorise separately.
**Hydrogen chloride, . Hydrogen needs one and chlorine needs one. One shared pair — a single bond — leaves hydrogen with a duplet and chlorine with an octet.
Water, . Oxygen (2, 6) needs two electrons, and each hydrogen can supply one. So oxygen forms two single bonds, one to each hydrogen. Oxygen ends with an octet — two bond pairs and two lone pairs — and each hydrogen with a duplet.
Ammonia, . Nitrogen (2, 5) needs three, so it forms three single bonds to three hydrogens. Nitrogen ends with an octet — three bond pairs and one lone pair.
Methane, . Carbon (2, 4) needs four, so it forms four single bonds to four hydrogens. Carbon ends with an octet and no lone pairs.
Carbon tetrachloride, . Carbon forms four single bonds, one to each chlorine. Carbon reaches an octet, and each chlorine reaches an octet with three lone pairs of its own.
Water, ammonia and methane are the same idea with different shortfalls. Oxygen is short two and takes two hydrogens; nitrogen short three takes three; carbon short four takes four. So the formula , , is read straight off the valence-electron count, and drawing the structure is a matter of arranging bond pairs and then filling in whatever lone pairs are left over.
Counting the electrons in a structure. For water, oxygen brings 6 valence electrons and the two hydrogens bring 1 each, giving 8 in the picture: 2 bond pairs (4 electrons) and 2 lone pairs (4 electrons). If your drawing has more or fewer electrons than the atoms brought, it is wrong** — and this check catches the usual error of adding a lone pair that has nowhere to come from.
How do ionic and covalent compounds differ in their properties?
Single, double and triple bonds first.
- Single bond — 1 shared pair: , , , , , ,
- Double bond — 2 shared pairs: , (), ethene
- Triple bond — 3 shared pairs: , ethyne
Carbon dioxide has two double bonds, not one. Carbon is short four electrons and each oxygen is short two, so carbon forms a double bond to each oxygen — — and every atom reaches an octet.
Now the comparison. Almost every difference below traces back to one structural fact: an ionic compound is a giant lattice held by strong electrostatic forces, and a covalent compound is made of discrete molecules held to each other only weakly.
- Structure — electrovalent: a giant lattice of ions; covalent: discrete molecules
- Melting and boiling point — electrovalent: high; covalent: low
- Physical state — electrovalent: hard crystalline solids; covalent: often gases or liquids, or soft solids
- Solubility in water — electrovalent: generally soluble; covalent: generally insoluble
- Solubility in organic solvents — electrovalent: generally insoluble; covalent: soluble in benzene, ether or
- Conduction as a solid — electrovalent: does not conduct; covalent: does not conduct
- Conduction when molten or dissolved — electrovalent: conducts; covalent: does not conduct
- Speed of reaction — electrovalent: fast; covalent: slow
Melting point. To melt an ionic solid you must overcome the strong electrostatic attraction throughout an entire lattice. To melt a covalent solid you need only separate the molecules from one another, and the forces between molecules are weak — the strong covalent bonds inside each molecule do not have to break at all. That is the whole reason for the difference, and it is the answer to "why does sodium chloride melt so much higher than water?"
Electrical conduction. Conduction requires charged particles free to move.
- A solid ionic compound has ions, but they are locked in fixed positions in the lattice, so it does not conduct
- Molten or dissolved, the lattice is broken up and the ions are mobile, so it does conduct
- A covalent compound has no ions at all, so it does not conduct in any state
So the ions must exist and be mobile — both conditions. Saying an ionic compound conducts because it contains ions is only half an answer, and the solid case is the marker's favourite way of testing whether you know the other half.
The exception worth knowing. Some polar covalent compounds ionise in water, even though they contain no ions to begin with. Hydrogen chloride is a covalent gas, and its solution in water conducts because it has broken into and . So "covalent compounds do not conduct" needs the qualifier "unless they ionise in water", and the acids of the next chapter are all in this category.
Speed of reaction. An ionic reaction in solution is fast because the ions are already present and only have to meet. A covalent reaction is slow because existing bonds have to be broken before anything can be made — and that is what the last section is about.
- Single bond — 1 shared pair: , , , , , ,
- Double bond — 2 shared pairs: , (), ethene
- Triple bond — 3 shared pairs: , ethyne
Carbon dioxide has two double bonds, not one. Carbon is short four electrons and each oxygen is short two, so carbon forms a double bond to each oxygen — — and every atom reaches an octet.
Now the comparison. Almost every difference below traces back to one structural fact: an ionic compound is a giant lattice held by strong electrostatic forces, and a covalent compound is made of discrete molecules held to each other only weakly.
- Structure — electrovalent: a giant lattice of ions; covalent: discrete molecules
- Melting and boiling point — electrovalent: high; covalent: low
- Physical state — electrovalent: hard crystalline solids; covalent: often gases or liquids, or soft solids
- Solubility in water — electrovalent: generally soluble; covalent: generally insoluble
- Solubility in organic solvents — electrovalent: generally insoluble; covalent: soluble in benzene, ether or
- Conduction as a solid — electrovalent: does not conduct; covalent: does not conduct
- Conduction when molten or dissolved — electrovalent: conducts; covalent: does not conduct
- Speed of reaction — electrovalent: fast; covalent: slow
Melting point. To melt an ionic solid you must overcome the strong electrostatic attraction throughout an entire lattice. To melt a covalent solid you need only separate the molecules from one another, and the forces between molecules are weak — the strong covalent bonds inside each molecule do not have to break at all. That is the whole reason for the difference, and it is the answer to "why does sodium chloride melt so much higher than water?"
Electrical conduction. Conduction requires charged particles free to move.
- A solid ionic compound has ions, but they are locked in fixed positions in the lattice, so it does not conduct
- Molten or dissolved, the lattice is broken up and the ions are mobile, so it does conduct
- A covalent compound has no ions at all, so it does not conduct in any state
So the ions must exist and be mobile — both conditions. Saying an ionic compound conducts because it contains ions is only half an answer, and the solid case is the marker's favourite way of testing whether you know the other half.
The exception worth knowing. Some polar covalent compounds ionise in water, even though they contain no ions to begin with. Hydrogen chloride is a covalent gas, and its solution in water conducts because it has broken into and . So "covalent compounds do not conduct" needs the qualifier "unless they ionise in water", and the acids of the next chapter are all in this category.
Speed of reaction. An ionic reaction in solution is fast because the ions are already present and only have to meet. A covalent reaction is slow because existing bonds have to be broken before anything can be made — and that is what the last section is about.
What happens when a covalent bond breaks?
A shared pair belongs to two atoms. When the bond breaks, the pair has to be divided — and there are exactly two ways to divide two electrons.
Either one atom takes both, or they take one each. The two routes give completely different products.
Heterolytic fission — one atom takes both electrons. The more electronegative atom keeps the whole shared pair. It thereby gains an extra electron and becomes negatively charged, while the other atom is left short of one and becomes positively charged.
The products are a pair of ions:
Chlorine is more electronegative than hydrogen, so chlorine takes both electrons of the shared pair and leaves as ; hydrogen leaves as , a bare proton.
That is the reason hydrogen chloride solution conducts. A covalent molecule has produced ions by splitting its bond unevenly — the exception noted in the comparison table, now explained.
Homolytic fission — one electron each. The shared pair is split evenly, so each atom leaves with one electron of the pair. Neither atom gains or loses charge, so the products are electrically neutral — but each carries an unpaired electron.
A neutral atom or group with an unpaired electron is a free radical:
Homolytic fission needs energy supplied as heat or light, and it happens most readily where the two atoms are identical — since neither has any advantage in the tug of war.
Free radicals are extremely reactive. An unpaired electron is exactly what an atom is trying to avoid; a free radical will attack almost anything to pair it up. That is why they are short-lived intermediates rather than substances you can bottle, and why a reaction that proceeds through them is so hard to stop once started.
Radical and free radical are two different things. The earlier part of this chapter defined a radical as a charged group of atoms that acts as a single unit — sulphate , ammonium , nitrate . A free radical is uncharged and has an unpaired electron.
- Charge — a radical is charged; a free radical is neutral
- Unpaired electron — a radical has none; a free radical has one
- How it is produced — a radical by the dissociation of a compound into ions; a free radical by homolytic fission
- Examples — radicals and ; free radicals and
The similarity of the names is the whole trap, and a question naming one when it means the other is testing exactly this distinction. Heterolytic fission gives ions and homolytic fission gives free radicals — remember the pairing and both definitions follow.
Why this matters for the rest of chemistry. Nearly every reaction is a bond broken and a bond made, and which of the two fissions occurs decides what kind of reaction it is. Reactions in water generally go through ions, which is why they are fast. Reactions in sunlight or at high temperature often go through free radicals, which is why they proceed in chains.
Either one atom takes both, or they take one each. The two routes give completely different products.
Heterolytic fission — one atom takes both electrons. The more electronegative atom keeps the whole shared pair. It thereby gains an extra electron and becomes negatively charged, while the other atom is left short of one and becomes positively charged.
The products are a pair of ions:
Chlorine is more electronegative than hydrogen, so chlorine takes both electrons of the shared pair and leaves as ; hydrogen leaves as , a bare proton.
That is the reason hydrogen chloride solution conducts. A covalent molecule has produced ions by splitting its bond unevenly — the exception noted in the comparison table, now explained.
Homolytic fission — one electron each. The shared pair is split evenly, so each atom leaves with one electron of the pair. Neither atom gains or loses charge, so the products are electrically neutral — but each carries an unpaired electron.
A neutral atom or group with an unpaired electron is a free radical:
Homolytic fission needs energy supplied as heat or light, and it happens most readily where the two atoms are identical — since neither has any advantage in the tug of war.
Free radicals are extremely reactive. An unpaired electron is exactly what an atom is trying to avoid; a free radical will attack almost anything to pair it up. That is why they are short-lived intermediates rather than substances you can bottle, and why a reaction that proceeds through them is so hard to stop once started.
Radical and free radical are two different things. The earlier part of this chapter defined a radical as a charged group of atoms that acts as a single unit — sulphate , ammonium , nitrate . A free radical is uncharged and has an unpaired electron.
- Charge — a radical is charged; a free radical is neutral
- Unpaired electron — a radical has none; a free radical has one
- How it is produced — a radical by the dissociation of a compound into ions; a free radical by homolytic fission
- Examples — radicals and ; free radicals and
The similarity of the names is the whole trap, and a question naming one when it means the other is testing exactly this distinction. Heterolytic fission gives ions and homolytic fission gives free radicals — remember the pairing and both definitions follow.
Why this matters for the rest of chemistry. Nearly every reaction is a bond broken and a bond made, and which of the two fissions occurs decides what kind of reaction it is. Reactions in water generally go through ions, which is why they are fast. Reactions in sunlight or at high temperature often go through free radicals, which is why they proceed in chains.
Exam tip
Exam tip: count the shortfall, then draw
Define electrovalent bonding as complete TRANSFER of electrons from metal to non-metal, both reaching an octet, the ions then held by electrostatic attraction. Define covalent bonding as mutual SHARING of pairs, each atom contributing equally.
Derive the formula from the electron count. Magnesium loses 2 and each chlorine takes 1, so . Calcium loses 2 and oxygen takes 2, so .
For an ionic structure, draw both ions separately, mark the transferred electron with an arrow, and write the charges.
For a covalent structure, count the shortfall: short 1 gives a single bond, short 2 a double, short 3 a triple. Then check the total electrons drawn equals what the atoms brought.
Do not forget the lone pairs — three on each chlorine in , two on oxygen in , one on nitrogen in , none on carbon in .
Carbon dioxide has TWO double bonds, .
Explain high melting point structurally: an ionic lattice needs strong electrostatic forces overcome throughout, while a covalent solid needs only weak intermolecular forces broken.
Conduction needs ions that are FREE TO MOVE — so a solid ionic compound does not conduct, and the same compound molten or dissolved does.
Name the exception: polar covalent ionises in water and its solution conducts.
And keep the fissions apart — heterolytic gives ions, homolytic gives free radicals — and a radical is charged while a free radical is neutral with an unpaired electron.
Derive the formula from the electron count. Magnesium loses 2 and each chlorine takes 1, so . Calcium loses 2 and oxygen takes 2, so .
For an ionic structure, draw both ions separately, mark the transferred electron with an arrow, and write the charges.
For a covalent structure, count the shortfall: short 1 gives a single bond, short 2 a double, short 3 a triple. Then check the total electrons drawn equals what the atoms brought.
Do not forget the lone pairs — three on each chlorine in , two on oxygen in , one on nitrogen in , none on carbon in .
Carbon dioxide has TWO double bonds, .
Explain high melting point structurally: an ionic lattice needs strong electrostatic forces overcome throughout, while a covalent solid needs only weak intermolecular forces broken.
Conduction needs ions that are FREE TO MOVE — so a solid ionic compound does not conduct, and the same compound molten or dissolved does.
Name the exception: polar covalent ionises in water and its solution conducts.
And keep the fissions apart — heterolytic gives ions, homolytic gives free radicals — and a radical is charged while a free radical is neutral with an unpaired electron.
Did you know
Why table salt survives a furnace and water boils in a kettle
There is a pair of substances in every kitchen that makes the ionic-covalent difference impossible to miss.
Water boils on a stove. Salt does not melt in an oven, a flame, or anything else a kitchen can produce — you would need a furnace.
Both are made of small, light atoms. Both are held together by strong bonds. The reason they behave so differently is not the strength of the bonds but what has to break when you heat them.
To boil water you do not break a single bond. Every water molecule leaves the liquid completely intact and floats off as a molecule of steam. All you have to overcome is the weak attraction between one molecule and the next. The covalent bonds inside each molecule are barely involved, and they are still there in the steam.
To melt salt you have no such shortcut. There is no molecule of sodium chloride to send off intact — a salt crystal is one continuous lattice in which every sodium ion is surrounded by chloride ions and every chloride by sodiums, throughout the whole grain. Melting means loosening every one of those attractions at once.
So the covalent substance gets to keep its strong bonds and break only weak ones; the ionic substance has only strong ones to break.
This is why so many covalent compounds are gases and liquids at room temperature while nearly every ionic compound is a solid. And it explains something otherwise odd about diamond: it is covalent, and it is one of the hardest and highest-melting substances there is. Diamond has no separate molecules either — it is a giant covalent lattice, so melting it means breaking actual covalent bonds.
The property that matters is not ionic versus covalent but lattice versus molecule. Salt and diamond are both lattices and both refuse to melt easily; water and carbon dioxide are both molecular and both leave as intact molecules when warmed. The ionic-covalent labels are a very good guide to melting point only because ionic compounds are always lattices and covalent ones usually are not.
Water boils on a stove. Salt does not melt in an oven, a flame, or anything else a kitchen can produce — you would need a furnace.
Both are made of small, light atoms. Both are held together by strong bonds. The reason they behave so differently is not the strength of the bonds but what has to break when you heat them.
To boil water you do not break a single bond. Every water molecule leaves the liquid completely intact and floats off as a molecule of steam. All you have to overcome is the weak attraction between one molecule and the next. The covalent bonds inside each molecule are barely involved, and they are still there in the steam.
To melt salt you have no such shortcut. There is no molecule of sodium chloride to send off intact — a salt crystal is one continuous lattice in which every sodium ion is surrounded by chloride ions and every chloride by sodiums, throughout the whole grain. Melting means loosening every one of those attractions at once.
So the covalent substance gets to keep its strong bonds and break only weak ones; the ionic substance has only strong ones to break.
This is why so many covalent compounds are gases and liquids at room temperature while nearly every ionic compound is a solid. And it explains something otherwise odd about diamond: it is covalent, and it is one of the hardest and highest-melting substances there is. Diamond has no separate molecules either — it is a giant covalent lattice, so melting it means breaking actual covalent bonds.
The property that matters is not ionic versus covalent but lattice versus molecule. Salt and diamond are both lattices and both refuse to melt easily; water and carbon dioxide are both molecular and both leave as intact molecules when warmed. The ionic-covalent labels are a very good guide to melting point only because ionic compounds are always lattices and covalent ones usually are not.
Exam relevance
How does chemical bonding carry into JEE Main and NEET?
Because bonding is the chapter every later chapter depends on, and Class 11 devotes an entire unit to it.
This is the foundation for Class 11 Chemistry Chemical Bonding and Molecular Structure, examined in both JEE Main and NEET. Everything on this page reappears there with the reasoning supplied. The octet rule is stated and then its exceptions are examined — molecules where the central atom ends with fewer than eight electrons or more than eight — and questions asking you to identify a compound that disobeys the octet rule are a standard type.
The structures you draw here become Lewis structures with a formal procedure. Class 11 gives an algorithm for them, adds formal charge as a way of choosing between competing structures, and introduces resonance for molecules that no single structure describes. The electron-counting check used on this page — the drawing must contain exactly the electrons the atoms brought — is the first step of that procedure.
Shape becomes predictable. Class 11 introduces VSEPR theory, which uses the bond pairs and lone pairs counted here to predict the actual shape of a molecule. This is where the lone pairs stop being decoration: water has two and is bent, ammonia has one and is pyramidal, methane has none and is tetrahedral. **Shape questions on , and are extremely common in both JEE Main and NEET, and they are answered from the lone-pair count on this page.
Single, double and triple become bond order, bond length and bond energy. Class 11 makes the trend quantitative — the higher the bond order the shorter and the stronger** the bond — so the triple bond in is the shortest and strongest, and questions ask you to rank bonds by length or energy. The pair-counting done here is what produces the bond order.
Polarity gets a measure. The polar covalent bond in , mentioned here as the exception that conducts in water, is treated with electronegativity difference and dipole moment in Class 11, and hybridisation and molecular orbital theory follow.
Free radicals become a mechanism. Class 11 Hydrocarbons uses homolytic fission for the free-radical substitution of alkanes, and heterolytic fission for the ionic mechanisms of Class 12 Haloalkanes and Haloarenes. The two fissions distinguished on this page are the two mechanism types organic chemistry is built on, and carbocations, carbanions and free radicals are examined as intermediates.
For NEET, Chemical Bonding is among the most heavily examined chapters in Chemistry, and Biology needs it too — Biomolecules assumes covalent bonding, and the hydrogen bonding that holds the two strands of DNA together is bonding of exactly this kind. Ionic and covalent character questions, shape questions and bond-order rankings are all standard NEET types.
What the questions look like. For board work, expect define electrovalent and covalent bonding, **draw the orbit structure of , , , , , , or , give three differences between ionic and covalent compounds with reasons, explain why a solid ionic compound does not conduct, and distinguish a radical from a free radical. Structures must show the charges or the lone pairs. For JEE Main and NEET, expect VSEPR shapes, bond-order rankings, octet-rule exceptions, formal charge and dipole moment.
How board and competitive emphasis differ. A board paper rewards a correctly drawn structure and a structural reason for a property. A competitive paper assumes both and asks for a shape, a bond angle or a polarity comparison.
The single trap that costs the most marks. Saying that an ionic compound conducts electricity because it contains ions. It contains ions in the solid state too, and the solid does not conduct — the ions must also be free to move, which happens only when molten or dissolved. The defence is to say "free" every time you say "ions"**, because that one word is the difference between the full mark and half of it.
This is the foundation for Class 11 Chemistry Chemical Bonding and Molecular Structure, examined in both JEE Main and NEET. Everything on this page reappears there with the reasoning supplied. The octet rule is stated and then its exceptions are examined — molecules where the central atom ends with fewer than eight electrons or more than eight — and questions asking you to identify a compound that disobeys the octet rule are a standard type.
The structures you draw here become Lewis structures with a formal procedure. Class 11 gives an algorithm for them, adds formal charge as a way of choosing between competing structures, and introduces resonance for molecules that no single structure describes. The electron-counting check used on this page — the drawing must contain exactly the electrons the atoms brought — is the first step of that procedure.
Shape becomes predictable. Class 11 introduces VSEPR theory, which uses the bond pairs and lone pairs counted here to predict the actual shape of a molecule. This is where the lone pairs stop being decoration: water has two and is bent, ammonia has one and is pyramidal, methane has none and is tetrahedral. **Shape questions on , and are extremely common in both JEE Main and NEET, and they are answered from the lone-pair count on this page.
Single, double and triple become bond order, bond length and bond energy. Class 11 makes the trend quantitative — the higher the bond order the shorter and the stronger** the bond — so the triple bond in is the shortest and strongest, and questions ask you to rank bonds by length or energy. The pair-counting done here is what produces the bond order.
Polarity gets a measure. The polar covalent bond in , mentioned here as the exception that conducts in water, is treated with electronegativity difference and dipole moment in Class 11, and hybridisation and molecular orbital theory follow.
Free radicals become a mechanism. Class 11 Hydrocarbons uses homolytic fission for the free-radical substitution of alkanes, and heterolytic fission for the ionic mechanisms of Class 12 Haloalkanes and Haloarenes. The two fissions distinguished on this page are the two mechanism types organic chemistry is built on, and carbocations, carbanions and free radicals are examined as intermediates.
For NEET, Chemical Bonding is among the most heavily examined chapters in Chemistry, and Biology needs it too — Biomolecules assumes covalent bonding, and the hydrogen bonding that holds the two strands of DNA together is bonding of exactly this kind. Ionic and covalent character questions, shape questions and bond-order rankings are all standard NEET types.
What the questions look like. For board work, expect define electrovalent and covalent bonding, **draw the orbit structure of , , , , , , or , give three differences between ionic and covalent compounds with reasons, explain why a solid ionic compound does not conduct, and distinguish a radical from a free radical. Structures must show the charges or the lone pairs. For JEE Main and NEET, expect VSEPR shapes, bond-order rankings, octet-rule exceptions, formal charge and dipole moment.
How board and competitive emphasis differ. A board paper rewards a correctly drawn structure and a structural reason for a property. A competitive paper assumes both and asks for a shape, a bond angle or a polarity comparison.
The single trap that costs the most marks. Saying that an ionic compound conducts electricity because it contains ions. It contains ions in the solid state too, and the solid does not conduct — the ions must also be free to move, which happens only when molten or dissolved. The defence is to say "free" every time you say "ions"**, because that one word is the difference between the full mark and half of it.
Key takeaways
Electrovalent and covalent bonding: quick revision
- Electrovalent bond: complete transfer of electrons from a metal to a non-metal so both reach a stable octet; the oppositely charged ions are held by electrostatic attraction.
- NaCl: (2, 8, 1) (2, 8); (2, 8, 7) (2, 8, 8).
- **MgCl**: (2, 8, 2) loses 2 (2, 8); two chlorines each gain 1. The subscript is electron bookkeeping.
- CaO: (2, 8, 8, 2) (2, 8, 8); (2, 6) (2, 8). Two electrons given, two taken, so no subscripts.
- Ionic compounds have no molecules — a solid is a giant lattice, and the formula is a formula unit giving the ratio of ions.
- Covalent bond: mutual sharing of one or more pairs, each atom contributing equally, so each reaches an octet or a duplet.
- Count the shortfall to get the bond. and — one pair, single. — two pairs, double . — three pairs, triple .
- ** single, H duplet and Cl octet. two single bonds, O octet with 2 lone pairs. three single bonds, N octet with 1 lone pair. four single bonds, C octet, no lone pairs. four single bonds.
- has two double bonds**, .
- Check the electron total in any structure against what the atoms brought.
- Melting point: ionic high — a whole lattice of strong electrostatic forces must be overcome; covalent low — only weak intermolecular forces break, and the covalent bonds survive intact.
- State: ionic hard crystalline solids; covalent often gases, liquids or soft solids.
- Solubility: ionic soluble in water and not in organic solvents; covalent the reverse.
- Conduction needs ions FREE TO MOVE. Solid ionic — no; molten or dissolved ionic — yes; covalent — no, except polar covalent like , which ionises in water.
- Reaction speed: ionic fast (ions already present); covalent slow (bonds must break first).
- Heterolytic fission — the more electronegative atom takes both electrons, giving ions: .
- Homolytic fission — one electron each, giving neutral free radicals with an unpaired electron: , driven by heat or light.
- A radical is charged (, ); a free radical is neutral with an unpaired electron () and is highly reactive.
Cover the structures and try drawing , and from the valence electrons alone, lone pairs included — if all three come out right, the rest of this chapter is arithmetic.
- NaCl: (2, 8, 1) (2, 8); (2, 8, 7) (2, 8, 8).
- **MgCl**: (2, 8, 2) loses 2 (2, 8); two chlorines each gain 1. The subscript is electron bookkeeping.
- CaO: (2, 8, 8, 2) (2, 8, 8); (2, 6) (2, 8). Two electrons given, two taken, so no subscripts.
- Ionic compounds have no molecules — a solid is a giant lattice, and the formula is a formula unit giving the ratio of ions.
- Covalent bond: mutual sharing of one or more pairs, each atom contributing equally, so each reaches an octet or a duplet.
- Count the shortfall to get the bond. and — one pair, single. — two pairs, double . — three pairs, triple .
- ** single, H duplet and Cl octet. two single bonds, O octet with 2 lone pairs. three single bonds, N octet with 1 lone pair. four single bonds, C octet, no lone pairs. four single bonds.
- has two double bonds**, .
- Check the electron total in any structure against what the atoms brought.
- Melting point: ionic high — a whole lattice of strong electrostatic forces must be overcome; covalent low — only weak intermolecular forces break, and the covalent bonds survive intact.
- State: ionic hard crystalline solids; covalent often gases, liquids or soft solids.
- Solubility: ionic soluble in water and not in organic solvents; covalent the reverse.
- Conduction needs ions FREE TO MOVE. Solid ionic — no; molten or dissolved ionic — yes; covalent — no, except polar covalent like , which ionises in water.
- Reaction speed: ionic fast (ions already present); covalent slow (bonds must break first).
- Heterolytic fission — the more electronegative atom takes both electrons, giving ions: .
- Homolytic fission — one electron each, giving neutral free radicals with an unpaired electron: , driven by heat or light.
- A radical is charged (, ); a free radical is neutral with an unpaired electron () and is highly reactive.
Cover the structures and try drawing , and from the valence electrons alone, lone pairs included — if all three come out right, the rest of this chapter is arithmetic.