The Hardest Substance and a Pencil Lead Are the Same Element
Learn where carbon occurs free and combined, what allotropy means, how the structures of diamond and graphite explain their opposite properties, and how fullerene differs from both.
How can the hardest known substance and a soft pencil lead be identical?
Because diamond and graphite are both nothing but carbon atoms — the same element, with no other substance present in either.
Burn a diamond in oxygen and it gives carbon dioxide. Burn graphite and it gives carbon dioxide. Chemically they are indistinguishable.
What differs is entirely the arrangement of the atoms. Diamond locks each carbon rigidly to four neighbours in three dimensions; graphite arranges them in flat sheets that slide over one another. That single difference in geometry produces the hardest substance there is and a material soft enough to write with. This page covers the first part of the ICSE Class 8 Chemistry chapter on carbon.
Burn a diamond in oxygen and it gives carbon dioxide. Burn graphite and it gives carbon dioxide. Chemically they are indistinguishable.
What differs is entirely the arrangement of the atoms. Diamond locks each carbon rigidly to four neighbours in three dimensions; graphite arranges them in flat sheets that slide over one another. That single difference in geometry produces the hardest substance there is and a material soft enough to write with. This page covers the first part of the ICSE Class 8 Chemistry chapter on carbon.
Where does carbon occur, and what is allotropy?
In the free state — as diamond, graphite and coal, found in the Earth's crust.
In the combined state — almost everywhere:
- Carbon dioxide in air, at about to by volume — small but essential, since it is the raw material of photosynthesis
- Carbonates in rocks — limestone, marble, chalk and dolomite, all containing
- Hydrocarbons — petroleum, natural gas, methane
- Carbohydrates, fats and proteins — so all food, and every living and once-living thing
- Fossil fuels — coal, coke, petroleum, formed from the remains of ancient organisms
Carbon is the element on which the whole of life is built, which is why it has a branch of chemistry to itself.
Allotropy is the existence of an element in two or more different physical forms having the same chemical properties but different physical properties. Those forms are called allotropes.
The key words are same chemical and different physical. Every allotrope of carbon burns in oxygen to give carbon dioxide — the chemistry is identical. Their hardness, density, colour, lustre and conductivity are wildly different.
Carbon's allotropes fall into two groups:
- Crystalline — the atoms are arranged in a regular, repeating pattern: diamond, graphite, fullerene
- Amorphous — no regular arrangement: coal, coke, charcoal, lampblack, gas carbon
The amorphous forms are the subject of the next part; the three crystalline ones are dealt with here.
Carbon is not alone in this. Sulphur has rhombic and monoclinic allotropes; oxygen has ordinary oxygen and ozone ; phosphorus has red and white forms. Carbon is simply the most dramatic example, because no other element's allotropes differ so extremely in their properties.
In the combined state — almost everywhere:
- Carbon dioxide in air, at about to by volume — small but essential, since it is the raw material of photosynthesis
- Carbonates in rocks — limestone, marble, chalk and dolomite, all containing
- Hydrocarbons — petroleum, natural gas, methane
- Carbohydrates, fats and proteins — so all food, and every living and once-living thing
- Fossil fuels — coal, coke, petroleum, formed from the remains of ancient organisms
Carbon is the element on which the whole of life is built, which is why it has a branch of chemistry to itself.
Allotropy is the existence of an element in two or more different physical forms having the same chemical properties but different physical properties. Those forms are called allotropes.
The key words are same chemical and different physical. Every allotrope of carbon burns in oxygen to give carbon dioxide — the chemistry is identical. Their hardness, density, colour, lustre and conductivity are wildly different.
Carbon's allotropes fall into two groups:
- Crystalline — the atoms are arranged in a regular, repeating pattern: diamond, graphite, fullerene
- Amorphous — no regular arrangement: coal, coke, charcoal, lampblack, gas carbon
The amorphous forms are the subject of the next part; the three crystalline ones are dealt with here.
Carbon is not alone in this. Sulphur has rhombic and monoclinic allotropes; oxygen has ordinary oxygen and ozone ; phosphorus has red and white forms. Carbon is simply the most dramatic example, because no other element's allotropes differ so extremely in their properties.
How do the structures of diamond and graphite explain their properties?
Every property of each follows from how its atoms are joined.
Diamond's structure. Each carbon atom is bonded to four other carbon atoms, arranged tetrahedrally, forming one continuous rigid three-dimensional network. There are no layers, no gaps and no loose atoms — the whole crystal is a single giant molecule.
What follows from it:
- Hardest natural substance. Every atom is held in place from four directions at once, so nothing can be pushed out of position.
- Does not conduct electricity. All four of each atom's outer electrons are used in bonding, so no free electrons remain to carry a current.
- Very high melting point. Melting means breaking the entire network.
- Brilliant lustre and sparkle, from a high refractive index — light entering the crystal is bent and reflected repeatedly before it escapes.
- High density, about , because the atoms are packed tightly.
- Transparent and colourless when pure.
Graphite's structure. Each carbon atom is bonded to only three others, giving flat hexagonal sheets like layers of chicken wire. Within a sheet the bonding is strong, but the sheets are held to one another only weakly, so they can slide past each other. Each atom's fourth outer electron is not used in bonding and is free to move along the sheet.
What follows from it:
- Soft and slippery. The sheets slide, which is why graphite marks paper and works as a lubricant.
- Conducts electricity well. Those free electrons drift along the sheets — making graphite the one non-metal that conducts.
- Greyish-black with a metallic lustre, and opaque.
- Lower density, about , because the widely spaced sheets leave more empty room. So diamond is about
times as dense as graphite, from the same atoms.
- High melting point as well, since the sheets themselves are strongly bonded.
The single sentence worth memorising. Diamond uses all four outer electrons in bonding and has none spare; graphite uses three and leaves one free. Hardness and non-conduction come from the first; softness and conduction from the second.
A misconception to clear. Graphite is soft between its layers, not within them. Each sheet is extremely strong — which is why carbon fibre, built from similar sheets, is used where strength matters. Soft describes how easily the layers slide, not how weak the bonds are.
Diamond's structure. Each carbon atom is bonded to four other carbon atoms, arranged tetrahedrally, forming one continuous rigid three-dimensional network. There are no layers, no gaps and no loose atoms — the whole crystal is a single giant molecule.
What follows from it:
- Hardest natural substance. Every atom is held in place from four directions at once, so nothing can be pushed out of position.
- Does not conduct electricity. All four of each atom's outer electrons are used in bonding, so no free electrons remain to carry a current.
- Very high melting point. Melting means breaking the entire network.
- Brilliant lustre and sparkle, from a high refractive index — light entering the crystal is bent and reflected repeatedly before it escapes.
- High density, about , because the atoms are packed tightly.
- Transparent and colourless when pure.
Graphite's structure. Each carbon atom is bonded to only three others, giving flat hexagonal sheets like layers of chicken wire. Within a sheet the bonding is strong, but the sheets are held to one another only weakly, so they can slide past each other. Each atom's fourth outer electron is not used in bonding and is free to move along the sheet.
What follows from it:
- Soft and slippery. The sheets slide, which is why graphite marks paper and works as a lubricant.
- Conducts electricity well. Those free electrons drift along the sheets — making graphite the one non-metal that conducts.
- Greyish-black with a metallic lustre, and opaque.
- Lower density, about , because the widely spaced sheets leave more empty room. So diamond is about
times as dense as graphite, from the same atoms.
- High melting point as well, since the sheets themselves are strongly bonded.
The single sentence worth memorising. Diamond uses all four outer electrons in bonding and has none spare; graphite uses three and leaves one free. Hardness and non-conduction come from the first; softness and conduction from the second.
A misconception to clear. Graphite is soft between its layers, not within them. Each sheet is extremely strong — which is why carbon fibre, built from similar sheets, is used where strength matters. Soft describes how easily the layers slide, not how weak the bonds are.
How do diamond, graphite and fullerene compare?
Fullerene is the third crystalline allotrope, and it is neither a network nor a sheet but a closed cage.
The best-known form contains sixty carbon atoms arranged as a hollow sphere of linked hexagons and pentagons, written . Larger and tube-shaped versions also exist, the tubes being called nanotubes.
Setting the three side by side:
- Structure. Diamond: rigid three-dimensional network, each atom bonded to four. Graphite: flat hexagonal sheets, each atom bonded to three. Fullerene: hollow spherical cage of sixty atoms.
- Hardness. Diamond: hardest known. Graphite: soft and slippery. Fullerene: soft, and the molecules are separate.
- Electrical conduction. Diamond: none. Graphite: good. Fullerene: poor on its own, but it becomes a conductor when other atoms are added to it.
- Appearance. Diamond: transparent, brilliant. Graphite: greyish-black, metallic lustre. Fullerene: a dark solid.
- Density. Diamond about , graphite about , fullerene lower still, being hollow.
Their uses, each following from a property:
Diamond —
- Jewellery, for its lustre and sparkle
- Cutting, drilling and grinding, for its hardness: glass cutters, rock drills, and the saws used to cut marble
- Abrasives in polishing powders
- Dies for drawing very fine wire
Graphite —
- Pencil leads, mixed with clay to set the hardness; more clay gives a harder pencil
- Electrodes in dry cells and in electrolysis, for its conductivity
- Lubricant for machinery too hot for oil, since graphite does not burn or decompose easily
- Crucibles for melting metals, for its high melting point
- A moderator in nuclear reactors
Fullerene —
- Lubricants, since the spherical molecules can roll
- Nanotechnology and electronics research
- Catalysts, and as a possible carrier for delivering medicines
The pattern across all three. Nothing in those lists is arbitrary. A glass cutter needs hardness, so diamond. An electrode needs conduction, so graphite. Both come from the same element, and the choice is made purely on structure — which is the whole reason allotropy is worth a chapter.
The best-known form contains sixty carbon atoms arranged as a hollow sphere of linked hexagons and pentagons, written . Larger and tube-shaped versions also exist, the tubes being called nanotubes.
Setting the three side by side:
- Structure. Diamond: rigid three-dimensional network, each atom bonded to four. Graphite: flat hexagonal sheets, each atom bonded to three. Fullerene: hollow spherical cage of sixty atoms.
- Hardness. Diamond: hardest known. Graphite: soft and slippery. Fullerene: soft, and the molecules are separate.
- Electrical conduction. Diamond: none. Graphite: good. Fullerene: poor on its own, but it becomes a conductor when other atoms are added to it.
- Appearance. Diamond: transparent, brilliant. Graphite: greyish-black, metallic lustre. Fullerene: a dark solid.
- Density. Diamond about , graphite about , fullerene lower still, being hollow.
Their uses, each following from a property:
Diamond —
- Jewellery, for its lustre and sparkle
- Cutting, drilling and grinding, for its hardness: glass cutters, rock drills, and the saws used to cut marble
- Abrasives in polishing powders
- Dies for drawing very fine wire
Graphite —
- Pencil leads, mixed with clay to set the hardness; more clay gives a harder pencil
- Electrodes in dry cells and in electrolysis, for its conductivity
- Lubricant for machinery too hot for oil, since graphite does not burn or decompose easily
- Crucibles for melting metals, for its high melting point
- A moderator in nuclear reactors
Fullerene —
- Lubricants, since the spherical molecules can roll
- Nanotechnology and electronics research
- Catalysts, and as a possible carrier for delivering medicines
The pattern across all three. Nothing in those lists is arbitrary. A glass cutter needs hardness, so diamond. An electrode needs conduction, so graphite. Both come from the same element, and the choice is made purely on structure — which is the whole reason allotropy is worth a chapter.
Exam tip
Exam tip: give the structure before the property
Never state a property without its structural reason. Graphite conducts electricity is weak; graphite conducts because each carbon is bonded to only three others, leaving one free electron per atom to move along the layers is complete.
The two numbers to keep straight: diamond bonds to four, graphite to three. Every difference follows from that.
Define allotropy with both halves — same chemical properties, different physical properties. Dropping either half loses the mark.
Say diamond has no free electrons as the reason it does not conduct.
For graphite's softness, say the layers slide over one another, and note that the layers themselves are strongly bonded.
Name the groups: crystalline allotropes are diamond, graphite and fullerene; amorphous ones are coal, coke, charcoal and lampblack.
For uses, pair each with the property responsible — hardness for cutting, conductivity for electrodes, high melting point for crucibles.
Quote fullerene as , a hollow cage of sixty atoms.
And remember that all allotropes of carbon burn to give carbon dioxide. That identical chemistry is what makes them allotropes rather than different substances.
The two numbers to keep straight: diamond bonds to four, graphite to three. Every difference follows from that.
Define allotropy with both halves — same chemical properties, different physical properties. Dropping either half loses the mark.
Say diamond has no free electrons as the reason it does not conduct.
For graphite's softness, say the layers slide over one another, and note that the layers themselves are strongly bonded.
Name the groups: crystalline allotropes are diamond, graphite and fullerene; amorphous ones are coal, coke, charcoal and lampblack.
For uses, pair each with the property responsible — hardness for cutting, conductivity for electrodes, high melting point for crucibles.
Quote fullerene as , a hollow cage of sixty atoms.
And remember that all allotropes of carbon burn to give carbon dioxide. That identical chemistry is what makes them allotropes rather than different substances.
Did you know
Why is a harder pencil made with more clay, not more graphite?
A pencil marked H is harder than one marked B, and both are darker and softer than you might expect for something called lead. There is no lead in either.
A pencil core is graphite mixed with clay and baked. Graphite alone would be too soft and crumbly to hold a point, so the clay binds it and stiffens it. Adding more clay gives a harder core that wears down slowly and leaves a fainter mark; adding more graphite gives a softer core that sheds easily and marks darkly.
So the H and B grades are really a clay-to-graphite ratio, and nothing about the graphite itself changes between them.
What makes the pencil work at all is the layered structure. Press the core on paper and the weakly held sheets slide off in flakes and stick to the fibres. Rubbing with an eraser lifts those flakes away again — which is why pencil can be erased and ink cannot.
A pencil core is graphite mixed with clay and baked. Graphite alone would be too soft and crumbly to hold a point, so the clay binds it and stiffens it. Adding more clay gives a harder core that wears down slowly and leaves a fainter mark; adding more graphite gives a softer core that sheds easily and marks darkly.
So the H and B grades are really a clay-to-graphite ratio, and nothing about the graphite itself changes between them.
What makes the pencil work at all is the layered structure. Press the core on paper and the weakly held sheets slide off in flakes and stick to the fibres. Rubbing with an eraser lifts those flakes away again — which is why pencil can be erased and ink cannot.
Key takeaways
Carbon and its crystalline allotropes: quick revision
- Carbon occurs free as diamond, graphite and coal, and combined as in air (about to by volume), carbonates in limestone, marble, chalk and dolomite, hydrocarbons, and all food and living matter.
- Allotropy is an element existing in two or more forms with the same chemical and different physical properties.
- Crystalline allotropes: diamond, graphite, fullerene. Amorphous: coal, coke, charcoal, lampblack, gas carbon.
- Other elements show it too — sulphur, oxygen (with ozone) and phosphorus.
- Diamond: each carbon bonded to four others in a rigid three-dimensional network. So it is the hardest substance, does not conduct (no free electrons), has a very high melting point, a brilliant lustre, and a density of about .
- Graphite: each carbon bonded to three others in flat hexagonal sheets that slide, with one free electron per atom. So it is soft and slippery, conducts electricity — the one conducting non-metal — is greyish-black with a metallic lustre, and has a density of about .
- Diamond is therefore about times as dense as graphite, from the same atoms.
- Graphite's layers slide easily but the sheets themselves are strongly bonded.
- Fullerene is a hollow cage, , of sixty atoms; soft, a poor conductor alone, and lower in density.
- Uses — diamond: jewellery, glass cutters, rock drills, abrasives, wire-drawing dies. Graphite: pencil leads with clay, electrodes, high-temperature lubricant, crucibles, reactor moderator. Fullerene: lubricants, nanotechnology, catalysts.
- All allotropes burn in oxygen to give carbon dioxide, which is what makes them allotropes of one element.
Try explaining each property of diamond and graphite from its structure without looking — being able to derive them rather than recall them is what this chapter is really testing.
- Allotropy is an element existing in two or more forms with the same chemical and different physical properties.
- Crystalline allotropes: diamond, graphite, fullerene. Amorphous: coal, coke, charcoal, lampblack, gas carbon.
- Other elements show it too — sulphur, oxygen (with ozone) and phosphorus.
- Diamond: each carbon bonded to four others in a rigid three-dimensional network. So it is the hardest substance, does not conduct (no free electrons), has a very high melting point, a brilliant lustre, and a density of about .
- Graphite: each carbon bonded to three others in flat hexagonal sheets that slide, with one free electron per atom. So it is soft and slippery, conducts electricity — the one conducting non-metal — is greyish-black with a metallic lustre, and has a density of about .
- Diamond is therefore about times as dense as graphite, from the same atoms.
- Graphite's layers slide easily but the sheets themselves are strongly bonded.
- Fullerene is a hollow cage, , of sixty atoms; soft, a poor conductor alone, and lower in density.
- Uses — diamond: jewellery, glass cutters, rock drills, abrasives, wire-drawing dies. Graphite: pencil leads with clay, electrodes, high-temperature lubricant, crucibles, reactor moderator. Fullerene: lubricants, nanotechnology, catalysts.
- All allotropes burn in oxygen to give carbon dioxide, which is what makes them allotropes of one element.
Try explaining each property of diamond and graphite from its structure without looking — being able to derive them rather than recall them is what this chapter is really testing.