Two Chlorine Atoms Share Electrons Because Neither Can Take One Outright
See how atoms that pull on electrons equally end up sharing them, draw the dot structures of hydrogen, chlorine, oxygen, nitrogen, methane and carbon tetrachloride, find out why HCl, ammonia and water are polar, and compare covalent compounds with ionic ones property by property.
Why do two identical atoms bond when neither can take the other's electron?
A chlorine atom, , is one electron short of an octet. Next to a sodium atom that problem is easy to solve — sodium simply hands over its outer electron, as the previous part showed.
But chlorine gas is made of two chlorine atoms bonded together, and there is no sodium in sight. Both atoms want one more electron, and both pull on electrons with exactly the same strength. Neither can win an electron from the other. Yet chlorine does not exist as separate atoms; it exists as molecules.
The way out is to share. Each chlorine atom puts one electron into a pair held between the two nuclei. That pair belongs to both atoms at once, so each now counts eight electrons round itself, and both have an octet.
A bond formed by sharing a pair of electrons is a covalent bond, and it holds together most of the substances you meet in a day:
- The water you drink and the oxygen and nitrogen you breathe
- The methane that reaches many city kitchens through gas pipelines
- The sugar in tea and the wax of a candle
So the picture of bonding now has two ends. When one atom pulls far harder than the other, electrons are transferred and ions form. When the two pull equally, electrons are shared evenly. In between lies the case this chapter spends most care on — two different atoms sharing a pair unevenly, which gives a polar covalent bond.
The difference in pull is the electronegativity difference from the periodic properties chapter, and it decides which kind of bond forms. That link is the thread running through this whole part.
This page covers the second part of the ICSE Class 10 Chemistry chapter on chemical bonding: forming a covalent bond, dot structures of six non-polar molecules, polar covalent bonds in HCl, NH3 and H2O, and the contrast between electrovalent and covalent compounds.
But chlorine gas is made of two chlorine atoms bonded together, and there is no sodium in sight. Both atoms want one more electron, and both pull on electrons with exactly the same strength. Neither can win an electron from the other. Yet chlorine does not exist as separate atoms; it exists as molecules.
The way out is to share. Each chlorine atom puts one electron into a pair held between the two nuclei. That pair belongs to both atoms at once, so each now counts eight electrons round itself, and both have an octet.
A bond formed by sharing a pair of electrons is a covalent bond, and it holds together most of the substances you meet in a day:
- The water you drink and the oxygen and nitrogen you breathe
- The methane that reaches many city kitchens through gas pipelines
- The sugar in tea and the wax of a candle
So the picture of bonding now has two ends. When one atom pulls far harder than the other, electrons are transferred and ions form. When the two pull equally, electrons are shared evenly. In between lies the case this chapter spends most care on — two different atoms sharing a pair unevenly, which gives a polar covalent bond.
The difference in pull is the electronegativity difference from the periodic properties chapter, and it decides which kind of bond forms. That link is the thread running through this whole part.
This page covers the second part of the ICSE Class 10 Chemistry chapter on chemical bonding: forming a covalent bond, dot structures of six non-polar molecules, polar covalent bonds in HCl, NH3 and H2O, and the contrast between electrovalent and covalent compounds.
How does a covalent bond form by sharing electrons to reach a duplet or octet?
Two atoms each contribute electrons to one or more shared pairs, and every shared pair counts towards the duplet or octet of both atoms.
The definition. A covalent bond is the chemical bond formed between two atoms by the mutual sharing of one or more pairs of electrons, each atom contributing equally, so that both attain a stable electronic configuration.
Words you need:
- Bond pair or shared pair — a pair of electrons shared between two atoms
- Lone pair — a pair of outer electrons that belongs to one atom and is not shared
- Covalency — the number of electrons an atom contributes to sharing, which equals the number of covalent bonds it forms
Single, double and triple bonds.
- One shared pair is a single bond, drawn as one line:
- Two shared pairs make a double bond:
- Three shared pairs make a triple bond:
Covalency of the common non-metals, which is the number of electrons each needs:
- Hydrogen — needs one more for a duplet
- Chlorine — covalency
- Oxygen — covalency
- Nitrogen — covalency
- Carbon — covalency
These are the same numbers as the valencies of groups 14 to 17 from the periodic table — eight minus the valence electrons — which is why those elements form covalent bonds so readily.
Conditions for a covalent bond:
- Both atoms are non-metals, or one is hydrogen, each needing electrons rather than having spare ones
- Both have high ionisation potentials, so neither gives up electrons
- Their electronegativities are equal or close, so neither can pull an electron away completely
Worked example — the hydrogen molecule. Each hydrogen atom has one electron and needs two for the duplet of helium.
- Each atom contributes its single electron to one shared pair
- Both nuclei attract that pair, and each atom now counts two electrons round itself
- Electrons in the molecule: , all in one bond pair, with no lone pairs
Why the shared pair counts for both atoms. The pair sits between the two nuclei and is attracted by both. That shared attraction is the bond itself — pull the atoms apart and each nucleus loses its hold on half the pair.
An everyday way to see the difference from ionic bonding. A sodium chloride crystal dissolved in water fills the solution with ions, while sugar dissolved in water leaves only sugar molecules. Sugar is held together by covalent bonds, and a covalent molecule contains no ions at all — which, as the last section shows, is why sugar solution does not conduct electricity.
One boundary case. Sharing does not require the two atoms to be the same element. In methane, one carbon atom shares four pairs with four hydrogen atoms, contributing one electron to each pair. What makes a bond covalent is the sharing, not the identity of the atoms.
The definition. A covalent bond is the chemical bond formed between two atoms by the mutual sharing of one or more pairs of electrons, each atom contributing equally, so that both attain a stable electronic configuration.
Words you need:
- Bond pair or shared pair — a pair of electrons shared between two atoms
- Lone pair — a pair of outer electrons that belongs to one atom and is not shared
- Covalency — the number of electrons an atom contributes to sharing, which equals the number of covalent bonds it forms
Single, double and triple bonds.
- One shared pair is a single bond, drawn as one line:
- Two shared pairs make a double bond:
- Three shared pairs make a triple bond:
Covalency of the common non-metals, which is the number of electrons each needs:
- Hydrogen — needs one more for a duplet
- Chlorine — covalency
- Oxygen — covalency
- Nitrogen — covalency
- Carbon — covalency
These are the same numbers as the valencies of groups 14 to 17 from the periodic table — eight minus the valence electrons — which is why those elements form covalent bonds so readily.
Conditions for a covalent bond:
- Both atoms are non-metals, or one is hydrogen, each needing electrons rather than having spare ones
- Both have high ionisation potentials, so neither gives up electrons
- Their electronegativities are equal or close, so neither can pull an electron away completely
Worked example — the hydrogen molecule. Each hydrogen atom has one electron and needs two for the duplet of helium.
- Each atom contributes its single electron to one shared pair
- Both nuclei attract that pair, and each atom now counts two electrons round itself
- Electrons in the molecule: , all in one bond pair, with no lone pairs
Why the shared pair counts for both atoms. The pair sits between the two nuclei and is attracted by both. That shared attraction is the bond itself — pull the atoms apart and each nucleus loses its hold on half the pair.
An everyday way to see the difference from ionic bonding. A sodium chloride crystal dissolved in water fills the solution with ions, while sugar dissolved in water leaves only sugar molecules. Sugar is held together by covalent bonds, and a covalent molecule contains no ions at all — which, as the last section shows, is why sugar solution does not conduct electricity.
One boundary case. Sharing does not require the two atoms to be the same element. In methane, one carbon atom shares four pairs with four hydrogen atoms, contributing one electron to each pair. What makes a bond covalent is the sharing, not the identity of the atoms.
How do you draw electron dot structures for H2, Cl2, O2, N2, CH4 and CCl4?
Share as many pairs as each atom needs to complete its octet or duplet, place every remaining outer electron as a lone pair, and check that the total matches the outer electrons you started with.
The method, used for every molecule below:
- Write the outer electrons of each atom
- Decide how many each atom needs, which is its covalency
- Share that many pairs between the atoms
- Place the remaining electrons as lone pairs
- Check that every atom has an octet, or a duplet for hydrogen, and that the electron count adds up
In a drawing, one atom's electrons are shown as dots and the other's as crosses, so that every shared pair shows one dot and one cross.
**1. Hydrogen, .
- Outer electrons**:
- One shared pair, no lone pairs
- Structure: , each hydrogen with a duplet
**2. Chlorine, .
- Outer electrons**:
- Each chlorine needs one, so one shared pair
- Each chlorine keeps three lone pairs
- Check: , and each chlorine counts
- Structure:
**3. Oxygen, .
- Outer electrons**:
- Each oxygen needs two, so two shared pairs — a double bond
- Each oxygen keeps two lone pairs
- Check: , and each oxygen counts
- Structure:
**4. Nitrogen, .
- Outer electrons**:
- Each nitrogen needs three, so three shared pairs — a triple bond
- Each nitrogen keeps one lone pair
- Check: , and each nitrogen counts
- Structure:
**5. Methane, .
- Outer electrons**:
- Carbon needs four, and each hydrogen needs one, so four shared pairs, one to each hydrogen
- No lone pairs anywhere
- Check: ; carbon counts eight, each hydrogen two
- The four bonds point to the corners of a tetrahedron, as far apart as possible
**6. Carbon tetrachloride, .
- Outer electrons**:
- Four shared pairs, one between carbon and each chlorine
- Each chlorine keeps three lone pairs
- Check:
Why all six are non-polar molecules.
- **In , , and the two atoms are identical, so they pull equally and the shared pairs sit exactly in the middle
- In and the four bonds are arranged symmetrically, so any pull in one direction is balanced by the others
The boundary case in that last point. Chlorine is more electronegative than carbon, so each C–Cl bond on its own is slightly polar. But the four bonds point to the corners of a tetrahedron and their effects cancel exactly. A molecule can contain polar bonds and still be non-polar overall — the shape decides.
The common error. Drawing oxygen with a single bond leaves each oxygen atom with only seven electrons. Running the octet check on every atom catches it** — if any atom is short, another pair must be shared.
The method, used for every molecule below:
- Write the outer electrons of each atom
- Decide how many each atom needs, which is its covalency
- Share that many pairs between the atoms
- Place the remaining electrons as lone pairs
- Check that every atom has an octet, or a duplet for hydrogen, and that the electron count adds up
In a drawing, one atom's electrons are shown as dots and the other's as crosses, so that every shared pair shows one dot and one cross.
**1. Hydrogen, .
- Outer electrons**:
- One shared pair, no lone pairs
- Structure: , each hydrogen with a duplet
**2. Chlorine, .
- Outer electrons**:
- Each chlorine needs one, so one shared pair
- Each chlorine keeps three lone pairs
- Check: , and each chlorine counts
- Structure:
**3. Oxygen, .
- Outer electrons**:
- Each oxygen needs two, so two shared pairs — a double bond
- Each oxygen keeps two lone pairs
- Check: , and each oxygen counts
- Structure:
**4. Nitrogen, .
- Outer electrons**:
- Each nitrogen needs three, so three shared pairs — a triple bond
- Each nitrogen keeps one lone pair
- Check: , and each nitrogen counts
- Structure:
**5. Methane, .
- Outer electrons**:
- Carbon needs four, and each hydrogen needs one, so four shared pairs, one to each hydrogen
- No lone pairs anywhere
- Check: ; carbon counts eight, each hydrogen two
- The four bonds point to the corners of a tetrahedron, as far apart as possible
**6. Carbon tetrachloride, .
- Outer electrons**:
- Four shared pairs, one between carbon and each chlorine
- Each chlorine keeps three lone pairs
- Check:
Why all six are non-polar molecules.
- **In , , and the two atoms are identical, so they pull equally and the shared pairs sit exactly in the middle
- In and the four bonds are arranged symmetrically, so any pull in one direction is balanced by the others
The boundary case in that last point. Chlorine is more electronegative than carbon, so each C–Cl bond on its own is slightly polar. But the four bonds point to the corners of a tetrahedron and their effects cancel exactly. A molecule can contain polar bonds and still be non-polar overall — the shape decides.
The common error. Drawing oxygen with a single bond leaves each oxygen atom with only seven electrons. Running the octet check on every atom catches it** — if any atom is short, another pair must be shared.
Why are HCl, NH3 and H2O polar, and how are their structures drawn?
In each molecule the shared pairs are pulled closer to the more electronegative atom, which becomes slightly negative while the hydrogen atoms become slightly positive.
A polar covalent bond is a covalent bond between two different atoms in which the shared pair is displaced towards the more electronegative one. No electron is transferred, so there are no full charges — only partial charges, written and .
The more electronegative atom in all three molecules is not hydrogen. Chlorine, nitrogen and oxygen each attract a shared pair more strongly than hydrogen does, so hydrogen always carries the end.
1. Hydrogen chloride, HCl.
- Outer electrons:
- One shared pair between hydrogen and chlorine; three lone pairs on chlorine
- Hydrogen has a duplet, chlorine an octet
- Polarity:
2. Ammonia, NH3.
- Outer electrons:
- Three shared pairs, one to each hydrogen; one lone pair on nitrogen
- Check: round nitrogen
- **Nitrogen carries **, each hydrogen
- The lone pair pushes the three bonds down, giving a pyramidal molecule
3. Water, H2O.
- Outer electrons:
- Two shared pairs, one to each hydrogen; two lone pairs on oxygen
- Check: round oxygen
- **Oxygen carries **, each hydrogen
- The two lone pairs bend the molecule, so water is V-shaped rather than straight
Why these molecules stay polar while carbon tetrachloride does not. In ammonia and water, the lone pairs force the bonds to one side, so the partial charges do not cancel — the molecule has a negative end and a positive end. Carbon tetrachloride has no lone pair on carbon and a fully symmetrical shape, so its bond polarities cancel.
Comparing bond polarity. Oxygen is more electronegative than nitrogen, so an O–H bond is more polar than an N–H bond. The larger the electronegativity difference, the more unevenly the pair is shared.
An everyday observation that shows polarity. Run a plastic comb through dry hair and bring it close to a thin, steady stream of water from a tap. The stream bends towards the comb. Water molecules have charged ends that respond to the charged comb, which is also why water is so good at pulling apart the ions of common salt.
The boundary case to state precisely. A polar covalent compound is still covalent. Pure hydrogen chloride gas consists of HCl molecules, not ions, and does not conduct electricity. Only when it dissolves in water does it form ions, as water molecules pull the proton away:
That reaction, and the lone pair on water's oxygen that makes it possible, are the subject of the next part on coordinate bonds.
A polar covalent bond is a covalent bond between two different atoms in which the shared pair is displaced towards the more electronegative one. No electron is transferred, so there are no full charges — only partial charges, written and .
The more electronegative atom in all three molecules is not hydrogen. Chlorine, nitrogen and oxygen each attract a shared pair more strongly than hydrogen does, so hydrogen always carries the end.
1. Hydrogen chloride, HCl.
- Outer electrons:
- One shared pair between hydrogen and chlorine; three lone pairs on chlorine
- Hydrogen has a duplet, chlorine an octet
- Polarity:
2. Ammonia, NH3.
- Outer electrons:
- Three shared pairs, one to each hydrogen; one lone pair on nitrogen
- Check: round nitrogen
- **Nitrogen carries **, each hydrogen
- The lone pair pushes the three bonds down, giving a pyramidal molecule
3. Water, H2O.
- Outer electrons:
- Two shared pairs, one to each hydrogen; two lone pairs on oxygen
- Check: round oxygen
- **Oxygen carries **, each hydrogen
- The two lone pairs bend the molecule, so water is V-shaped rather than straight
Why these molecules stay polar while carbon tetrachloride does not. In ammonia and water, the lone pairs force the bonds to one side, so the partial charges do not cancel — the molecule has a negative end and a positive end. Carbon tetrachloride has no lone pair on carbon and a fully symmetrical shape, so its bond polarities cancel.
Comparing bond polarity. Oxygen is more electronegative than nitrogen, so an O–H bond is more polar than an N–H bond. The larger the electronegativity difference, the more unevenly the pair is shared.
An everyday observation that shows polarity. Run a plastic comb through dry hair and bring it close to a thin, steady stream of water from a tap. The stream bends towards the comb. Water molecules have charged ends that respond to the charged comb, which is also why water is so good at pulling apart the ions of common salt.
The boundary case to state precisely. A polar covalent compound is still covalent. Pure hydrogen chloride gas consists of HCl molecules, not ions, and does not conduct electricity. Only when it dissolves in water does it form ions, as water molecules pull the proton away:
That reaction, and the lone pair on water's oxygen that makes it possible, are the subject of the next part on coordinate bonds.
How do electrovalent and covalent compounds differ in state, melting point and conductivity?
Electrovalent compounds are built of ions held in a lattice, while covalent compounds are built of separate molecules with only weak forces between them — and almost every difference in their properties follows from that.
1. Constituent particles.
- Electrovalent: positive and negative ions
- Covalent: neutral molecules
2. State at room temperature.
- Electrovalent: hard crystalline solids, such as sodium chloride
- Covalent: gases such as methane and ammonia, liquids such as water and carbon tetrachloride, or soft solids such as wax, sugar and naphthalene
3. Melting and boiling points.
- Electrovalent: high, because strong attractions between ions in every direction must be overcome
- Covalent: low, because only the weak forces between molecules must be overcome
The boundary case inside that point. The covalent bonds inside a molecule are strong. Melting ice or boiling water separates whole water molecules from one another; it does not break a single O–H bond. Low melting points reflect weak forces between molecules, not weak covalent bonds.
4. Conductivity of electricity.
- Electrovalent: do not conduct as solids; conduct when molten or in aqueous solution, because the ions become free to move
- Covalent: do not conduct in any state, because they contain no ions — except polar covalent compounds such as HCl, which ionise in water and then conduct
5. Ionisation or dissociation in solution.
- Electrovalent: dissociate — the ions already exist and simply separate
- Polar covalent: ionise — ions are formed for the first time by reaction with water
- Non-polar covalent: neither
6. Solubility.
- Electrovalent: generally soluble in water, insoluble in organic solvents such as kerosene
- Covalent: generally insoluble in water but soluble in organic solvents — with exceptions such as sugar, alcohol, ammonia and hydrogen chloride, which dissolve in water
7. Speed of reactions.
- Electrovalent: reactions in solution are fast, since the ions only need to meet
- Covalent: reactions are usually slower, since bonds inside molecules must be broken
Worked comparison — three solutions in a circuit with a bulb.
- Common salt solution — the bulb glows. Sodium chloride has dissociated into mobile ions
- Sugar solution — the bulb does not glow. Sugar dissolves as molecules, with no ions
- Hydrochloric acid — the bulb glows. Hydrogen chloride is covalent but has ionised in water
Two of the three substances are covalent, yet one of them conducts — the case that shows why the conductivity rule for covalent compounds needs its exception.
An everyday check of the melting-point rule. A candle's wax melts from the gentle warmth a short distance below the flame, while grains of common salt left on a hot tawa stay solid and unchanged. Wax is a covalent compound of large molecules; salt is an ionic lattice.
One more boundary case worth knowing. Diamond and graphite are pure carbon joined entirely by covalent bonds, yet they have extremely high melting points, and graphite even conducts electricity. Their atoms are not in separate small molecules but in one continuous network, so melting would mean breaking covalent bonds themselves. The low-melting rule applies to molecular covalent substances, which are what this chapter describes.
1. Constituent particles.
- Electrovalent: positive and negative ions
- Covalent: neutral molecules
2. State at room temperature.
- Electrovalent: hard crystalline solids, such as sodium chloride
- Covalent: gases such as methane and ammonia, liquids such as water and carbon tetrachloride, or soft solids such as wax, sugar and naphthalene
3. Melting and boiling points.
- Electrovalent: high, because strong attractions between ions in every direction must be overcome
- Covalent: low, because only the weak forces between molecules must be overcome
The boundary case inside that point. The covalent bonds inside a molecule are strong. Melting ice or boiling water separates whole water molecules from one another; it does not break a single O–H bond. Low melting points reflect weak forces between molecules, not weak covalent bonds.
4. Conductivity of electricity.
- Electrovalent: do not conduct as solids; conduct when molten or in aqueous solution, because the ions become free to move
- Covalent: do not conduct in any state, because they contain no ions — except polar covalent compounds such as HCl, which ionise in water and then conduct
5. Ionisation or dissociation in solution.
- Electrovalent: dissociate — the ions already exist and simply separate
- Polar covalent: ionise — ions are formed for the first time by reaction with water
- Non-polar covalent: neither
6. Solubility.
- Electrovalent: generally soluble in water, insoluble in organic solvents such as kerosene
- Covalent: generally insoluble in water but soluble in organic solvents — with exceptions such as sugar, alcohol, ammonia and hydrogen chloride, which dissolve in water
7. Speed of reactions.
- Electrovalent: reactions in solution are fast, since the ions only need to meet
- Covalent: reactions are usually slower, since bonds inside molecules must be broken
Worked comparison — three solutions in a circuit with a bulb.
- Common salt solution — the bulb glows. Sodium chloride has dissociated into mobile ions
- Sugar solution — the bulb does not glow. Sugar dissolves as molecules, with no ions
- Hydrochloric acid — the bulb glows. Hydrogen chloride is covalent but has ionised in water
Two of the three substances are covalent, yet one of them conducts — the case that shows why the conductivity rule for covalent compounds needs its exception.
An everyday check of the melting-point rule. A candle's wax melts from the gentle warmth a short distance below the flame, while grains of common salt left on a hot tawa stay solid and unchanged. Wax is a covalent compound of large molecules; salt is an ionic lattice.
One more boundary case worth knowing. Diamond and graphite are pure carbon joined entirely by covalent bonds, yet they have extremely high melting points, and graphite even conducts electricity. Their atoms are not in separate small molecules but in one continuous network, so melting would mean breaking covalent bonds themselves. The low-melting rule applies to molecular covalent substances, which are what this chapter describes.
Exam tip
What earns full marks in a covalent bonding answer?
Show the outer electrons of each atom, the shared pairs, the lone pairs, and a count that proves every octet is complete.
- Start with the electronic configuration of each atom and state how many electrons it needs
- Use dots for one atom and crosses for the other, so each shared pair shows who contributed what
- Draw every lone pair — examiners mark missing lone pairs on chlorine, oxygen and nitrogen
- Name the bond type: single in , double in , triple in
- Check octets atom by atom, and duplets on hydrogen
- Add up the electrons in the structure and compare with the total outer electrons
- **Mark and on polar molecules, with the negative end on the more electronegative atom
- Say why a molecule is polar or non-polar — electronegativity difference plus shape
- Give a reason with every property in a comparison: ions or molecules, strong or weak forces
- Mention the HCl exception when stating that covalent compounds do not conduct
The misconception to name. Covalent compounds do not have low melting points because covalent bonds are weak. Covalent bonds are strong; the forces between molecules are weak**, and it is only those that melting overcomes. Writing covalent bonds are weak as the reason loses the mark.
A second trap. Calling carbon tetrachloride polar because chlorine is more electronegative than carbon. Each bond is polar, but the symmetrical shape cancels them, so the molecule is non-polar — a distinction that appears directly in questions.
- Start with the electronic configuration of each atom and state how many electrons it needs
- Use dots for one atom and crosses for the other, so each shared pair shows who contributed what
- Draw every lone pair — examiners mark missing lone pairs on chlorine, oxygen and nitrogen
- Name the bond type: single in , double in , triple in
- Check octets atom by atom, and duplets on hydrogen
- Add up the electrons in the structure and compare with the total outer electrons
- **Mark and on polar molecules, with the negative end on the more electronegative atom
- Say why a molecule is polar or non-polar — electronegativity difference plus shape
- Give a reason with every property in a comparison: ions or molecules, strong or weak forces
- Mention the HCl exception when stating that covalent compounds do not conduct
The misconception to name. Covalent compounds do not have low melting points because covalent bonds are weak. Covalent bonds are strong; the forces between molecules are weak**, and it is only those that melting overcomes. Writing covalent bonds are weak as the reason loses the mark.
A second trap. Calling carbon tetrachloride polar because chlorine is more electronegative than carbon. Each bond is polar, but the symmetrical shape cancels them, so the molecule is non-polar — a distinction that appears directly in questions.
Did you know
Why does the nitrogen in the air hardly react with anything?
Air is roughly four-fifths nitrogen, and every breath carries it into your lungs and straight back out, completely unchanged. Oxygen from the same breath is used by your body; the nitrogen might as well not be there. Food packets are often filled with nitrogen for the same reason — it sits beside the food without reacting, keeping out the oxygen that would make oil in snacks go stale.
The explanation is the dot structure you drew in this lesson.
- **Each nitrogen atom, , needs three electrons
- The two atoms share three pairs — a triple bond**,
- Six electrons packed between two small nuclei hold the atoms together extremely tightly
Breaking a triple bond takes a great deal of energy, far more than breaking the single bond in chlorine or the double bond in oxygen. So at ordinary temperatures, almost nothing can pull a nitrogen molecule apart.
Yet plants need nitrogen to build proteins, and they cannot use the gas directly. Nature has two ways of breaking that bond.
- Lightning supplies enough energy in a flash for nitrogen and oxygen to combine: . The products dissolve in rain and reach the soil
- Bacteria living in the root nodules of pulses — gram, peas, beans and groundnut — can break the bond using special enzymes, which is why farmers grow pulses to enrich the soil
Compare chlorine, whose atoms share just one pair. Chlorine reacts with a great many substances, attacking metals and bleaching dyes, because its single bond is far easier to break. Two gases, one needing one shared pair and the other needing three, and the number of shared pairs goes a long way towards explaining why one is dangerously reactive and the other almost inert.
And the same count explains a detail of the nitrogen molecule's electrons. With three pairs shared, each nitrogen atom has only one lone pair left. Every electron is accounted for: in bonds and in lone pairs make the outer electrons the two atoms brought.
The explanation is the dot structure you drew in this lesson.
- **Each nitrogen atom, , needs three electrons
- The two atoms share three pairs — a triple bond**,
- Six electrons packed between two small nuclei hold the atoms together extremely tightly
Breaking a triple bond takes a great deal of energy, far more than breaking the single bond in chlorine or the double bond in oxygen. So at ordinary temperatures, almost nothing can pull a nitrogen molecule apart.
Yet plants need nitrogen to build proteins, and they cannot use the gas directly. Nature has two ways of breaking that bond.
- Lightning supplies enough energy in a flash for nitrogen and oxygen to combine: . The products dissolve in rain and reach the soil
- Bacteria living in the root nodules of pulses — gram, peas, beans and groundnut — can break the bond using special enzymes, which is why farmers grow pulses to enrich the soil
Compare chlorine, whose atoms share just one pair. Chlorine reacts with a great many substances, attacking metals and bleaching dyes, because its single bond is far easier to break. Two gases, one needing one shared pair and the other needing three, and the number of shared pairs goes a long way towards explaining why one is dangerously reactive and the other almost inert.
And the same count explains a detail of the nitrogen molecule's electrons. With three pairs shared, each nitrogen atom has only one lone pair left. Every electron is accounted for: in bonds and in lone pairs make the outer electrons the two atoms brought.
Exam relevance
How does covalent bonding prepare you for JEE and NEET Chemistry?
This is foundation work for Class 11 Chemical Bonding and Molecular Structure, a chapter examined in both JEE Main and NEET Chemistry and used throughout organic chemistry afterwards.
Where dot structures lead. Class 11 calls them Lewis structures and adds formal charge, used to choose the best structure for a molecule or ion. The discipline you practise here — counting outer electrons, sharing pairs and checking every octet — is exactly the procedure, applied to larger molecules such as carbon dioxide, ozone and nitric acid, and to molecules that break the octet rule.
Where shapes lead. The pyramidal ammonia and bent water molecules you meet here are explained in Class 11 by VSEPR theory, in which lone pairs repel more strongly than bond pairs. Predicting shapes and bond angles from the number of bond pairs and lone pairs is a standard question type in both exams, and it begins with the lone-pair counts you write in this lesson.
Where polarity leads. Class 11 measures polarity as the dipole moment. A molecule with polar bonds but a symmetrical shape has zero dipole moment — the carbon tetrachloride case here — and comparing dipole moments of molecules such as water, ammonia, carbon dioxide and boron trifluoride is a recurring objective question.
Where the bond count leads. Class 11 introduces bond order, bond length and bond enthalpy, and later molecular orbital theory. The single, double and triple bonds of chlorine, oxygen and nitrogen are the standard examples: higher bond order means a shorter, stronger bond, which is the explanation for nitrogen's inertness.
Where the property comparison leads. The distinction between forces within molecules and forces between them becomes the study of hydrogen bonding and intermolecular forces, which explains why water boils so much higher than similar molecules.
Question types to expect. At this level: dot structures, polar versus non-polar, and property comparisons with reasons. In competitive papers: Lewis structures with formal charge, VSEPR shapes and bond angles, dipole moment comparisons, bond order, and assertion-reason items linking structure to properties.
The single trap that costs marks. Equating polar bonds with a polar molecule. Carbon tetrachloride and carbon dioxide both contain polar bonds and both have zero dipole moment, because their shapes are symmetrical. NEET and JEE Main both set options that only this distinction separates.
A second trap. Forgetting lone pairs when predicting shape. Water has four electron pairs round oxygen but only two bonds, so it is bent, not straight — and a candidate who counts only bonds gets both the shape and the polarity wrong.
Board versus competitive emphasis. The ICSE paper marks a correct dot diagram with all lone pairs and a reasoned property table; a competitive paper marks a shape, an angle or an order of dipole moments. The transferable habit is counting every outer electron before drawing anything, because every later bonding model is built on that count.
Where dot structures lead. Class 11 calls them Lewis structures and adds formal charge, used to choose the best structure for a molecule or ion. The discipline you practise here — counting outer electrons, sharing pairs and checking every octet — is exactly the procedure, applied to larger molecules such as carbon dioxide, ozone and nitric acid, and to molecules that break the octet rule.
Where shapes lead. The pyramidal ammonia and bent water molecules you meet here are explained in Class 11 by VSEPR theory, in which lone pairs repel more strongly than bond pairs. Predicting shapes and bond angles from the number of bond pairs and lone pairs is a standard question type in both exams, and it begins with the lone-pair counts you write in this lesson.
Where polarity leads. Class 11 measures polarity as the dipole moment. A molecule with polar bonds but a symmetrical shape has zero dipole moment — the carbon tetrachloride case here — and comparing dipole moments of molecules such as water, ammonia, carbon dioxide and boron trifluoride is a recurring objective question.
Where the bond count leads. Class 11 introduces bond order, bond length and bond enthalpy, and later molecular orbital theory. The single, double and triple bonds of chlorine, oxygen and nitrogen are the standard examples: higher bond order means a shorter, stronger bond, which is the explanation for nitrogen's inertness.
Where the property comparison leads. The distinction between forces within molecules and forces between them becomes the study of hydrogen bonding and intermolecular forces, which explains why water boils so much higher than similar molecules.
Question types to expect. At this level: dot structures, polar versus non-polar, and property comparisons with reasons. In competitive papers: Lewis structures with formal charge, VSEPR shapes and bond angles, dipole moment comparisons, bond order, and assertion-reason items linking structure to properties.
The single trap that costs marks. Equating polar bonds with a polar molecule. Carbon tetrachloride and carbon dioxide both contain polar bonds and both have zero dipole moment, because their shapes are symmetrical. NEET and JEE Main both set options that only this distinction separates.
A second trap. Forgetting lone pairs when predicting shape. Water has four electron pairs round oxygen but only two bonds, so it is bent, not straight — and a candidate who counts only bonds gets both the shape and the polarity wrong.
Board versus competitive emphasis. The ICSE paper marks a correct dot diagram with all lone pairs and a reasoned property table; a competitive paper marks a shape, an angle or an order of dipole moments. The transferable habit is counting every outer electron before drawing anything, because every later bonding model is built on that count.
Key takeaways
What must you be able to do from this part?
One way of bonding, six non-polar structures, three polar ones and a property comparison.
- A covalent bond is formed by the mutual sharing of one or more pairs of electrons, each atom contributing equally, so both reach a duplet or octet
- Bond pairs are shared; lone pairs are not; covalency is the number of electrons an atom shares
- Covalencies: hydrogen , chlorine , oxygen , nitrogen , carbon
- One, two and three shared pairs make single, double and triple bonds
- Covalent bonds form between non-metals with high ionisation potentials and equal or similar electronegativities
- ****: one shared pair, electrons
- ****: one shared pair and three lone pairs on each atom, electrons
- ****: a double bond and two lone pairs on each atom, electrons
- ****: a triple bond and one lone pair on each atom, electrons
- ****: four shared pairs and no lone pairs, electrons
- ****: four shared pairs and three lone pairs on each chlorine, electrons
- These six are non-polar — identical atoms, or a symmetrical shape that cancels polar bonds
- A polar covalent bond shares its pair unevenly, giving partial charges and but no ions
- HCl: one shared pair, three lone pairs on chlorine,
- **: three shared pairs and one lone pair on nitrogen, pyramidal
- : two shared pairs and two lone pairs on oxygen, V-shaped
- Electrovalent compounds: ions, crystalline solids, high melting and boiling points, conduct when molten or dissolved, dissociate, soluble in water, fast reactions
- Covalent compounds: molecules, gases, liquids or soft solids, low melting and boiling points, generally non-conducting, soluble in organic solvents, slower reactions
- HCl is covalent but ionises in water and then conducts
- Low melting points come from weak forces between molecules**, not from weak covalent bonds
The sharpest self-test is nine molecules and a blank page. Draw every structure from this lesson with all its lone pairs, count the electrons in each, and say whether each molecule is polar — then check whether your count and your answer agree.
- A covalent bond is formed by the mutual sharing of one or more pairs of electrons, each atom contributing equally, so both reach a duplet or octet
- Bond pairs are shared; lone pairs are not; covalency is the number of electrons an atom shares
- Covalencies: hydrogen , chlorine , oxygen , nitrogen , carbon
- One, two and three shared pairs make single, double and triple bonds
- Covalent bonds form between non-metals with high ionisation potentials and equal or similar electronegativities
- ****: one shared pair, electrons
- ****: one shared pair and three lone pairs on each atom, electrons
- ****: a double bond and two lone pairs on each atom, electrons
- ****: a triple bond and one lone pair on each atom, electrons
- ****: four shared pairs and no lone pairs, electrons
- ****: four shared pairs and three lone pairs on each chlorine, electrons
- These six are non-polar — identical atoms, or a symmetrical shape that cancels polar bonds
- A polar covalent bond shares its pair unevenly, giving partial charges and but no ions
- HCl: one shared pair, three lone pairs on chlorine,
- **: three shared pairs and one lone pair on nitrogen, pyramidal
- : two shared pairs and two lone pairs on oxygen, V-shaped
- Electrovalent compounds: ions, crystalline solids, high melting and boiling points, conduct when molten or dissolved, dissociate, soluble in water, fast reactions
- Covalent compounds: molecules, gases, liquids or soft solids, low melting and boiling points, generally non-conducting, soluble in organic solvents, slower reactions
- HCl is covalent but ionises in water and then conducts
- Low melting points come from weak forces between molecules**, not from weak covalent bonds
The sharpest self-test is nine molecules and a blank page. Draw every structure from this lesson with all its lone pairs, count the electrons in each, and say whether each molecule is polar — then check whether your count and your answer agree.