2.3.1 Carbon bonding and the variety of organic compounds

Every carbon atom forms four covalent bonds

1. Carbon has atomic number $6$ and the electron arrangement 2.4, so it has four outer electrons.
   1. It shares all four to gain a full outer shell, so each carbon atom forms four covalent bonds.
   2. In methane, ${CH}_4$, one carbon atom is bonded to four hydrogen atoms.
2. These bonds can be to hydrogen, oxygen, nitrogen, the halogens or other carbon atoms.
3. Carbon can also form double bonds, as in carbon dioxide and ethene, where two shared pairs join the same two atoms.

Carbon atoms form long chains, branches and rings

1. Bonds between carbon atoms are strong, so carbon atoms can link into chains of almost any length.
2. Chains can be straight or branched.
3. Carbon atoms can also join into rings, such as the ring of six carbon atoms in cyclohexane.
4. Poly(ethene) contains chains of thousands of carbon atoms, while methane has only one.
5. Chains, branches and rings can be combined in so many ways that there are millions of carbon compounds.

Carbon forms families of compounds with similar reactions

1. Members of a family of carbon compounds share the same general structure and differ mainly in the length of the carbon chain.
   1. Methane, ethane and propane belong to the alkanes, where each member has one more ${CH}_2$ than the one before.
2. Members of a family react in similar ways, because they contain the same group of atoms that takes part in reactions.
   1. Ethanol and propanol are alcohols, and each contains an $OH$ group.
3. Knowing how one member reacts lets chemists predict how the others will react.

Carbon bonding underlies natural and synthetic organic compounds

1. Living things are built from natural organic compounds, such as sugars, fats, proteins and DNA.
   1. Starch and cellulose in plants are long chains of sugar units joined together.
2. Chemists make synthetic organic compounds, including plastics, medicines, dyes and fuel additives.
3. Many synthetic compounds are made from crude oil, which is a mixture of hydrocarbons.
4. Natural and synthetic compounds rely on the same features of carbon, namely four bonds, chains and rings.

2.3.2 Diamond, graphite, fullerenes and graphene

Diamond, graphite, graphene and fullerenes are all carbon

1. All four are made only of carbon atoms joined by covalent bonds.
2. They have very different properties because the atoms are bonded differently and arranged in different shapes.
3. Diamond, graphite and graphene are giant covalent structures.
4. Fullerenes are molecules, such as buckminsterfullerene, $C_{60}$, which is shaped like a hollow football.

Four covalent bonds per atom make diamond hard

1. In diamond, each carbon atom is covalently bonded to four others in a rigid three-dimensional network.
2. Melting diamond means breaking many strong covalent bonds, which needs a huge amount of energy, so its melting point is very high.
3. The rigid network makes diamond extremely hard, which is why it is used on the tips of cutting tools and drills.
4. Diamond does not conduct electricity, because all four outer electrons of each atom are in bonds, leaving no delocalised electrons.

Layers and delocalised electrons explain graphite's properties

1. In graphite, each carbon atom forms three covalent bonds, making flat layers of hexagonal rings.
2. The fourth outer electron from each atom is delocalised, so graphite conducts electricity along the layers.
3. The layers are held together only by weak forces, so they slide over each other easily.
   1. This makes graphite soft and slippery, which is why it is used in pencils and as a lubricant.
4. Graphite still has a very high melting point, because melting breaks the strong covalent bonds within the layers.

Graphene sheets and hollow fullerene molecules

1. Graphene is one atom thick, so it is very light and almost transparent.
2. Its strong covalent bonds make it extremely strong for its mass.
3. Each carbon atom in graphene has a delocalised electron, so graphene conducts electricity very well and could be used in flexible electronics.
4. Fullerenes such as $C_{60}$ are separate molecules, so the weak intermolecular forces between them give much lower melting points than diamond.
5. Carbon nanotubes are fullerenes shaped as cylinders, which are very strong and conduct electricity, so they can reinforce materials.
6. The hollow shape of fullerenes lets them carry drug molecules into the body.

2.3.3 Bond strength, changes of state and predicting states

Changing state means overcoming forces between particles

1. Melting and boiling need energy to be transferred to the substance to overcome the attractions holding its particles together.
2. The stronger the attractions, the more energy is needed and the higher the melting and boiling points.
3. Freezing and condensing release energy as the attractions form again.
4. What has to be overcome depends on the type of structure, which is why melting points range from $-210{}^{\circ }C$ for nitrogen to over $3500{}^{\circ }C$ for diamond.

Simple molecular substances have low melting and boiling points

1. Inside each molecule, the atoms are held by strong covalent bonds.
2. Between the molecules there are only weak intermolecular forces.
   1. Melting or boiling separates the molecules, so only the intermolecular forces are overcome and the covalent bonds stay intact.
3. Little energy is needed, so these substances are often gases or liquids at room temperature.
4. Larger molecules have stronger intermolecular forces, so they have higher boiling points.
   1. This is why pentane boils at $36{}^{\circ }C$, while methane boils at $-162{}^{\circ }C$.

Giant structures have high melting points

1. In an ionic lattice, melting means overcoming strong electrostatic attractions between huge numbers of oppositely charged ions.
   1. Magnesium oxide, with $2+$ and $2-$ ions, has stronger attractions than sodium chloride and melts at about $2850{}^{\circ }C$.
2. In a giant covalent structure, melting breaks many strong covalent bonds, so diamond and silicon dioxide melt at very high temperatures.
3. In a metal, the strong attraction between positive ions and delocalised electrons means most metals need a lot of energy to melt.
4. Boiling points of giant structures are even higher, because the particles must be separated completely.

Predicting state from melting and boiling points

1. Below its melting point, a substance is a solid.
2. Between its melting and boiling points, it is a liquid.
3. Above its boiling point, it is a gas.
4. Take care with negative values, since $-50{}^{\circ }C$ is colder than $-20{}^{\circ }C$.
5. At its melting point, a substance can be present as both solid and liquid.

2.3.4 Bulk properties and types of bonding

Bulk properties come from bonding and structure

1. A single atom or molecule has no melting point and cannot conduct, because these are properties of the whole material.
2. Bulk properties depend on the type of bond, how strong it is compared with any intermolecular forces, and how the bonds are arranged.
3. Electrical conductivity needs charged particles that are free to move, either ions or delocalised electrons.
4. Hardness and brittleness depend on whether layers of particles can move past each other without the bonding breaking.

Hard, brittle ionic compounds conduct when molten or dissolved

1. Strong attractions between ions give ionic compounds high melting points.
2. Solid ionic compounds do not conduct, because their ions are fixed in the lattice.
   1. When molten or dissolved in water, the ions can move, so the compound conducts electricity.
3. Ionic crystals are brittle, because a sharp blow shifts a layer so that ions with the same charge line up and repel.
4. Many ionic compounds, such as sodium chloride, dissolve in water.

Simple molecular substances melt easily and do not conduct

1. Only weak intermolecular forces act between the molecules, so melting and boiling points are low.
2. Simple molecular substances do not conduct electricity, because the molecules have no overall charge and there are no free electrons.
3. Their solids are usually soft, because little force is needed to push the molecules apart.
4. Many do not dissolve in water, although some, such as sugar, do.

Giant covalent structures are hard, while polymers are flexible

1. Giant covalent substances such as diamond and silicon dioxide have very high melting points, because many strong covalent bonds must break.
2. They are hard, because the rigid network of bonds resists being pushed out of shape.
3. Most do not conduct electricity, because all their outer electrons are in bonds, with graphite the exception.
4. In a polymer, strong covalent bonds run along each chain, while weaker intermolecular forces hold neighbouring chains together.
   1. The chains can slide past each other, so many polymers are flexible and soften when heated.
   2. Polymers melt at higher temperatures than small molecules, because their long chains have stronger intermolecular forces.

Delocalised electrons and layers of ions explain metal properties

1. Metals conduct electricity, because their delocalised electrons can move through the structure.
2. Metals also conduct heat well, because the delocalised electrons carry energy through the metal.
3. Metals are malleable and ductile, because layers of ions can slide over each other without the metallic bonding breaking.
4. Most metals have high melting points, because of the strong attraction between the ions and the delocalised electrons.
5. In an alloy, atoms of different sizes disrupt the layers, so the layers slide less easily and the alloy is harder.

2.3.5 Nanoparticles: scale, surface area, uses and risks

Nanoparticles are between 1 and 100 nanometres across

1. One nanometre (nm) is $1\times 10^{-9}m$.
2. A typical atom is about $0.2nm$ across, so a $1nm$ nanoparticle is only about five atoms wide.
   1. A $100nm$ nanoparticle is about 500 atoms wide.
3. Nanoparticles are therefore much larger than simple molecules but far smaller than grains of dust or sand.

Smaller particles have larger surface area to volume ratios

1. For a cube of side $L$, the surface area is $6L^2$ and the volume is $L^3$, so the ratio is $\frac{6}{L}$.
2. When the side of a cube decreases by a factor of 10, its surface area to volume ratio increases by a factor of 10.
3. A nanoparticle therefore has a large fraction of its atoms at the surface compared with a lump of the same material.
4. Surface atoms are the ones that react or act as a catalyst, so nanoparticles often behave differently from the bulk material.

Large surface area gives nanoparticles their useful properties

1. Less material is needed for the same effect, because so much of each nanoparticle is surface.
2. Nanoparticles of platinum and other metals make efficient catalysts, for example in car catalytic converters.
3. Silver nanoparticles kill bacteria, so they are added to wound dressings, socks and some food containers.
4. Titanium dioxide and zinc oxide nanoparticles in sun creams absorb ultraviolet light but are too small to leave a white layer on the skin.
5. Nanoparticles are also being developed to deliver drugs to particular cells and for use in electronics and sensors.

The risks of nanoparticles are not fully understood

1. Nanoparticles are small enough to be breathed in deeply and may pass into the bloodstream and cells.
2. Their large surface area makes them more reactive, so they could be toxic in ways the bulk material is not.
3. Silver nanoparticles washed out of clothing could harm useful bacteria in soil and water.
4. Nanoparticle products are fairly new, so their long-term effects on health and the environment are still being studied.
5. This uncertainty is why nanoparticle products need thorough testing and clear labelling before sale.