S2.4 From models to materials

Scope. The bonding triangle is used for binary compounds only, with electronegativities from the data booklet; percentage ionic character is read from the triangle, never calculated. Examples of alloys (brass, bronze, steel and others) illustrate the ideas and need not be memorized. Monomer structures are given or deduced from the polymer, never recalled. Condensation polymers (polyamides and polyesters), and the formation and hydrolysis of biological macromolecules in general terms, are HL material; the detailed structures of starch, proteins and other biochemicals are not studied.

Bonding as a continuum

Three models of bonding

Chemists describe the strong bonds that hold substances together with three models. In the ionic model, electrons are transferred from a metal to a non-metal and the resulting oppositely charged ions attract one another in a lattice. In the covalent model, atoms of non-metals share pairs of electrons. In the metallic model, metal atoms release their valence electrons into a "sea" of delocalized electrons that holds a lattice of cations together.

Each model treats bonding as one pure type, yet very few substances are described perfectly by any one of them.

Why real bonds lie between the models

The deciding quantity is electronegativity: the ability of an atom to attract a shared pair of electrons towards itself. When two bonded atoms have the same electronegativity, as in ${{\mathrm{Cl}{\vphantom{X}}_{\smash[t]{2}}}}$, the bonding electrons are shared equally and the bond is purely covalent. When the electronegativities differ a little, as in ${{\mathrm{HCl}}}$, the shared pair is pulled towards the more electronegative atom. Chlorine carries a partial negative charge, δ−, and hydrogen a partial positive charge, δ+, so the bond is polar covalent and has some ionic character.

As the difference grows, the shared pair is pulled further towards one atom until, in practice, it belongs to that atom alone and the bond is best described as ionic. Even then the transfer is not complete. A small cation with a high charge, such as ${{\mathrm{Al}{\vphantom{X}}^{3+}}}$, distorts (polarizes) the electron cloud of a large anion such as ${{\mathrm{Cl}{\vphantom{X}}^{-}}}$, drawing electron density back into the region between the ions. This gives the "ionic" bond some covalent character.

The bonding continuum is a good example of how scientific models work. The categories "ionic", "covalent" and "metallic" are simplifications. They remain useful because each predicts a typical set of properties, but they fail at the boundaries, where substances behave partly like one type and partly like another. Aluminium chloride, for example, is a compound of a metal and a non-metal, yet it sublimes at about 180 °C and its vapour contains ${{\mathrm{Al}{\vphantom{X}}_{\smash[t]{2}}\mathrm{Cl}{\vphantom{X}}_{\smash[t]{6}}}}$ molecules, which is behaviour expected of a covalent substance.

The bonding triangle

A line from covalent to ionic captures only part of the picture, because a bond between two elements of low electronegativity is neither ionic nor covalent but metallic. The bonding triangle, also called the van Arkel–Ketelaar triangle, adds this third type. It is printed in the data booklet.

The two axes

Each point on the triangle is fixed by two numbers worked out from the electronegativities $\chi_\text{a}$ and $\chi_\text{b}$ of the two elements (data booklet):

The electronegativity difference, $\Delta\chi$, measures how unequally the bonding electrons are shared, so it sets the ionic character: the higher a point lies, the more ionic the bond. The average electronegativity, $\Sigma\chi$, measures how strongly the two atoms together hold on to electrons. A low average means that both atoms release electrons easily, as metals do; a high average means that both hold electrons tightly, as non-metals do. The symbol $\Sigma$ normally means a sum, but on the triangle $\Sigma\chi$ is the average of the two values.

Corners, edges and regions

• The base ($\Delta\chi = 0$) holds the elements, whose bonds join identical atoms. Metals lie towards the left (caesium, one of the least electronegative elements, is at the left corner), metalloids such as silicon near the middle, and non-metals towards the right, ending with fluorine at the right corner.

• The top corner stands for the most ionic bonding, as in caesium fluoride, formed from one of the least and the most electronegative elements.

• The metallic region (bottom left) contains elements and compounds with a low average electronegativity and a small difference: metals and alloys.

• The covalent and polar covalent regions (bottom right and middle right) contain substances with a high average electronegativity. Low on this side the sharing is nearly equal (covalent); higher up it becomes unequal (polar covalent).

• The ionic region (top) contains substances with a large electronegativity difference.

The scale on the right of the triangle reads off the approximate percentage covalent and ionic character from the height of a point, for example 25 % covalent and 75 % ionic at $\Delta\chi \approx 2.3$, and 50 % of each at $\Delta\chi \approx 1.6$. The boundaries between the regions are drawn as lines, but the change in bonding across them is gradual.

The compound with the largest electronegativity difference has the most ionic bonding. Fluorides of reactive metals have the largest differences of all: sodium fluoride, for example, has $\Delta\chi = 3.1$ and $\Sigma\chi = 2.45$. Points such as this, and caesium fluoride itself ($\Delta\chi = 3.2$, $\Sigma\chi = 2.4$), fall just outside the printed triangle, to the right of its apex, because the printed figure is an idealized shape drawn from rounded values. Such compounds are classified as ionic, with more than 90 % ionic character.

Placing a binary compound on the triangle

1. Look up the electronegativity of each element in the data booklet.

2. Calculate the difference, $\Delta\chi$ (always positive).

3. Calculate the average, $\Sigma\chi$.

4. Plot the point: $\Sigma\chi$ across, $\Delta\chi$ up.

5. Read the region the point falls in, and read the scale on the right at the same height for the percentage covalent and ionic character.

The triangle with a range of substances plotted shows how the bonding type changes with both electronegativities. Sodium chloride and potassium bromide are ionic; carbon dioxide and hydrogen chloride are polar covalent; methane is close to purely covalent; and an alloy of magnesium and aluminium is metallic.

Predicting properties from the bonding triangle

Because each region corresponds to a type of bonding, the position of a substance predicts its typical properties. The bonding type alone does not fix the properties, however, because a covalent substance may be made of small molecules or of a giant network, and the two behave very differently.

Substances close to a boundary have intermediate properties. Aluminium oxide, for example, sits on the boundary between ionic and polar covalent: it has a high melting point like an ionic solid, but it is amphoteric, reacting with both acids and bases, and its bonding has considerable covalent character.

The oxides of period 3: a continuum across a period

Across period 3, from sodium to sulfur, the electronegativity of the element increases from 0.9 to 2.6 while that of oxygen stays at 3.4. The electronegativity difference in the oxides therefore falls steadily, and the oxides move down and to the right across the bonding triangle, from the ionic region into the polar covalent region.

The change in bonding produces a change in structure and in properties. The first three oxides are giant ionic lattices, silicon dioxide is a giant covalent network, and the last two are simple molecular substances. The acid–base character follows the same trend, from basic ionic oxides to acidic covalent oxides.

Electronegativity differences from data booklet values; melting points are rounded reference values.

Silicon dioxide shows why structure must be considered as well as bonding. Its point on the triangle lies near those of ${{\mathrm{P}{\vphantom{X}}_{\smash[t]{4}}\mathrm{O}{\vphantom{X}}_{\smash[t]{10}}}}$ and ${{\mathrm{SO}{\vphantom{X}}_{\smash[t]{3}}}}$, yet it melts at about 1700 °C while sulfur trioxide melts at 17 °C. Melting silicon dioxide breaks strong covalent bonds throughout a giant network; melting sulfur trioxide only overcomes weak forces between small molecules.

Limitations of the bonding triangle

• It applies only to elements and binary compounds. A compound containing a polyatomic ion, such as ${{\mathrm{NH}{\vphantom{X}}_{\smash[t]{4}}\mathrm{Br}}}$ or ${{\mathrm{NaOH}}}$, has more than one type of bond and cannot be placed with two electronegativities.

• Electronegativity values are themselves approximate, and an element does not have exactly the same electronegativity in every compound.

• It gives no information about structure: ${{\mathrm{SiO}{\vphantom{X}}_{\smash[t]{2}}}}$ and ${{\mathrm{CO}{\vphantom{X}}_{\smash[t]{2}}}}$ are both polar covalent, but one is a giant network and the other a gas of small molecules.

• Elements near the middle of the base are borderline. Tin ($\chi = 2.0$) falls just inside the covalent region; tin is a metal at room temperature, but below about 13 °C it slowly changes into grey tin, a brittle form with a diamond-like covalent structure.

• The boundaries are drawn as sharp lines, although real bonding changes gradually across them.

Composite materials

Engineers often combine materials with different bonding to get properties that no single material has. A composite contains two or more distinct materials that remain separate within the finished structure. Reinforced concrete is a familiar example. Concrete, made mainly of ionic and covalent compounds such as calcium silicates, is strong when squeezed (in compression) but cracks easily when stretched (in tension). Steel bars, held by metallic bonding, are strong in tension. Embedding the steel in the concrete gives a material strong in both, which is why it is used for beams and bridges. Fibreglass, made of glass fibres in a polymer resin, combines the stiffness of the fibres with the toughness of the plastic.

Alloys

Pure metals are often too soft, too weak or too easily corroded for practical use. Pure iron bends easily and rusts; pure gold scratches and bends so easily that most jewellery is made from gold alloys. Mixing a metal with small amounts of other elements produces materials with better properties.

How alloys form: non-directional metallic bonding

In a metal, the attraction between the cations and the delocalized electrons acts equally in all directions. The bonding is non-directional: it does not depend on any particular atom sitting next to any other. An atom of a different element can therefore take the place of a metal atom in the lattice, or squeeze into a gap between atoms, and still be held by the sea of electrons. This is why molten metals mix readily with one another and why so many alloys exist. Covalent bonds, by contrast, are directional, so a covalent network cannot accept foreign atoms in the same way.

There are two ways in which the second element fits into the lattice. In a substitutional alloy, such as brass (copper and zinc), atoms of similar size replace some of the original atoms. In an interstitial alloy, such as steel (iron and carbon), much smaller atoms occupy the spaces between the metal atoms.

Alloys are mixtures, not compounds

An alloy is a mixture rather than a compound. Its composition can vary continuously: brass may contain anywhere from about 5 % to 45 % zinc, and each composition is still brass. There is no fixed ratio of atoms and no chemical formula, and the elements are not joined by new ionic or covalent bonds; they share the same delocalized electrons. A compound, in contrast, always has the same fixed composition and properties that differ completely from those of its elements. Most alloys are homogeneous mixtures, with the same composition throughout, much like a solid solution.

Explaining the properties of alloys

The properties of an alloy follow from two facts: the metallic bonding is still there, and the lattice is no longer perfectly regular.

Harder and stronger, but less malleable

A pure metal is malleable because its layers of identical cations can slide over one another while the delocalized electrons keep holding them together. In an alloy, atoms of a different size distort the layers around them. The layers are no longer smooth planes, so it takes a much larger force to make them slide. The alloy is therefore harder and stronger than the pure metal, but less malleable and less ductile.

Electrical and thermal conductivity

Alloys conduct electricity because they still contain delocalized electrons that move when a potential difference is applied. They usually conduct less well than the pure metal: the irregular lattice scatters the moving electrons more. Brass, for example, conducts electricity roughly a quarter as well as pure copper, which is why electrical wiring is made of pure copper.

Resistance to corrosion

Some alloys resist corrosion far better than their main metal. Stainless steel contains at least about 10.5 % chromium (a common grade has 18 % chromium and 8 % nickel). The chromium forms a thin, tightly bound layer of chromium(III) oxide on the surface that stops oxygen and water reaching the iron, and the layer reforms if it is scratched.

Lower melting point

Some alloys melt at a lower temperature than any of their components. Traditional solder, about 63 % tin and 37 % lead, melts at 183 °C, below both tin (231.9 °C) and lead (327.5 °C, data booklet values). The low melting point lets it join electrical components without damaging them. Lead-free solders based on tin with small amounts of silver and copper have now largely replaced it because lead is toxic.

Some common alloys

Compositions are typical values; each alloy is made in many grades.

Steel: one element, many properties

Steel shows how strongly a small amount of a second element changes a metal. Pure iron is soft. Adding carbon, whose atoms are much smaller than iron atoms, places carbon atoms in the gaps of the iron lattice. Each one distorts the surrounding layers, so the more carbon, the harder and stronger the metal, and the less it can be bent or drawn into wire.

Mild steel, with roughly 0.1–0.3 % carbon, is still easy to shape, so it is pressed into car bodies and rolled into girders. High-carbon steel, with roughly 0.6–1 % carbon, is much harder and keeps a sharp edge, so it is used for cutting tools and springs, but it is more brittle. Cast iron, with 2–4 % carbon, is very hard but so brittle that it breaks rather than bends; it is cast into shape for engine blocks and cookware. Adding other metals gives further control: chromium and nickel give stainless steel, and manganese makes steel tough enough for railway lines.

Titanium alloys show the same idea used for a different purpose. Titanium is about 40 % less dense than steel and resists corrosion, and alloying it with aluminium and vanadium makes it far stronger. Ti-6Al-4V is the most widely used titanium alloy; titanium alloys make up about 15 % of the mass of a Boeing 787 aircraft, where strength for a given mass matters more than cost.

Polymers

Many of the most useful materials are made of enormous molecules. A single chain of poly(ethene) can contain tens of thousands of carbon atoms. Such molecules are built by joining very many small molecules, end to end, into a chain.

The monomer becomes part of the polymer; the section of the chain that each monomer contributes is called the repeating unit, the smallest part of the chain that repeats over and over. A polymer molecule is a macromolecule, but its formula is written using the repeating unit and a subscript $n$, the number of times the unit repeats. In a real sample $n$ varies from one chain to the next, so a polymer has an average chain length rather than one exact molar mass.

Natural and synthetic polymers

Polymers are not only manufactured. Living things build natural polymers from small molecules, and many of the materials used for thousands of years, such as wool, silk, cotton and wood, are made of them. Synthetic polymers are made industrially, mostly from compounds obtained from crude oil.

Synthetic polymers include poly(ethene), poly(propene), poly(chloroethene) (PVC), poly(phenylethene) (polystyrene) and poly(tetrafluoroethene) (PTFE), which are made from alkenes, and nylon, PET and Kevlar, which are made in a different way. The synthetic polymers that can be moulded into shape are commonly called plastics.

How the structure of plastics explains their properties

The properties of a plastic depend on three features of its structure: the strong covalent bonds within each chain, the length of the chains, and the forces between neighbouring chains. In poly(ethene) the chains are non-polar, so neighbouring chains attract only by London (dispersion) forces. Each force is weak, but long chains touch along their whole length and tangle around each other, so the total attraction is large. Longer chains therefore give higher melting points and greater strength.

Branching and cross-linking

Chains that are unbranched can line up and pack closely, increasing the London forces between them. Branches keep chains apart and weaken these forces. The two common forms of poly(ethene) show the effect clearly.

Values are typical ranges; they vary between grades.

Polar groups have a similar effect to closer packing. The C–Cl bonds of PVC are polar, so its chains attract each other by dipole–dipole forces as well as London forces, which makes pure PVC rigid enough for window frames and drainpipes. Adding small molecules called plasticizers between the chains pushes them apart and makes PVC flexible enough for cable insulation.

In some polymers, covalent bonds called cross-links join neighbouring chains. Heating cannot make the chains slide past one another without breaking these bonds, so the material does not soften and chars instead. Such polymers are called thermosets, in contrast to thermoplastics, which soften on heating and can be remoulded. The vulcanization of rubber with sulfur, which forms sulfur cross-links between chains, turns soft, sticky natural rubber into the hard-wearing rubber of tyres.

Plastic waste and biodegradability

The same features that make plastics useful make them a problem once thrown away. Microorganisms cannot break down the long hydrocarbon chains of poly(ethene), poly(propene) or polystyrene: the C–C and C–H bonds are strong and non-polar, and the chains contain no functional groups that the microorganisms' enzymes can attack. These plastics are not biodegradable, so they persist in the environment for decades or centuries. Sunlight and wear break them into smaller and smaller pieces; fragments less than 5 mm across, called microplastics, are now found in oceans, soils and living organisms.

About 460 million tonnes of plastics were used worldwide in 2019, and of the plastic waste produced, only about 9 % was recycled; roughly 19 % was incinerated, about half went to landfill, and the rest was dumped, burned in the open or leaked into the environment (OECD, 2022).

• Recycling needs plastics to be sorted by type, because different polymers do not mix well when melted. A resin identification code stamped on the item identifies the polymer.

• Incineration recovers energy but releases carbon dioxide; burning PVC also releases hydrogen chloride, ${{\mathrm{HCl}}}$, a toxic and acidic gas, so it needs gas scrubbing.

• Biodegradable plastics contain bonds that can be broken by hydrolysis, such as the ester links in polyesters. Polylactic acid (PLA), made from plant sugars, breaks down in industrial composting. Designing a greener polymer combines such features: a renewable feedstock, bonds that can be hydrolysed once the product is discarded, and a high atom economy.

Resin identification codes printed inside the recycling symbol on plastic items.

Addition polymers

Alkenes are the starting point for most of the plastics made in the largest amounts. Their carbon–carbon double bond is the reactive part. A double bond is made of two parts, and one of them is weaker than a single bond and breaks more easily. When this part breaks in each monomer, the two electrons it held are used to form new single bonds to neighbouring monomers, and a long saturated chain grows.

In addition polymerization the double bond of each monomer opens and becomes a single bond, and every monomer is linked to the next. Industrially it needs a catalyst or initiator, and often high pressure: the original process for poly(ethene) used pressures of over 1000 times atmospheric pressure. The equation for ethene is:

Atom economy of addition polymerization

All the atoms of every monomer end up in the polymer, and nothing else is formed. The mass of polymer equals the mass of monomer that reacted, so the atom economy of an addition polymerization is 100 %. This is one reason why addition polymers are cheap to make on an enormous scale. The percentage yield can still be below 100 % if some monomer does not react.

Which molecules can form addition polymers?

A molecule can act as a monomer for addition polymerization only if it contains a C=C double bond. Ethene, propene, chloroethene, phenylethene and tetrafluoroethene all qualify. Ethane, chloroethane or cyclohexane cannot form addition polymers, however reactive their other bonds may be, because they have no double bond to open.

Drawing the repeating unit of an addition polymer

1. Draw the monomer with its C=C double bond in the middle and all four groups attached to the two carbon atoms (above and below).

2. Replace the C=C with a C–C single bond. Keep every substituent on the carbon atom it was attached to.

3. Add a continuation bond to the outside of each carbon atom: these are the bonds that join the unit to its neighbours.

4. Draw square brackets so that the continuation bonds pass through them, and write $n$ after the closing bracket.

The repeating units of four common addition polymers follow this pattern. Only the two carbon atoms of the original double bond are in the main chain; every other group, however large, hangs from the chain as a side group.

Working backwards: the monomer from the polymer

To find the monomer of an addition polymer, find the smallest unit of two main-chain carbon atoms that repeats, remove the continuation bonds and put a double bond between the two carbon atoms. With the side groups kept in place, the result is the monomer.

Condensation polymers (HL)

Condensation reactions

Two condensation reactions matter most for polymers. A carboxylic acid reacts with an alcohol to form an ester, joined by an ester link, ${{{-}\mathrm{COO}{-}}}$. A carboxylic acid reacts with an amine to form an amide, joined by an amide link, ${{{-}\mathrm{CONH}{-}}}$. In each case the ${{{-}\mathrm{OH}}}$ of the acid and a hydrogen atom from the other molecule leave together as water:

In the laboratory, amides and polyamides are usually made from an acyl chloride, ${{\mathrm{R}{-}\mathrm{COCl}}}$, instead of a carboxylic acid, because it reacts faster and at room temperature. The small molecule released is then hydrogen chloride, ${{\mathrm{HCl}}}$, instead of water.

From condensation reaction to condensation polymer

A molecule with only one functional group can react only once: ethanoic acid and ethanol give a single ester molecule, and the reaction stops. To build a chain, each monomer needs two functional groups, one at each end, so that after reacting at one end it can still react at the other. There are two arrangements:

• two different monomers, each with two identical groups: a diol with a dicarboxylic acid (giving a polyester), or a diamine with a dicarboxylic acid or diacyl dichloride (giving a polyamide);

• one monomer with two different groups: a hydroxy acid such as lactic acid (2-hydroxypropanoic acid) gives a polyester, and an amino acid gives a polyamide.

A condensation polymer differs from an addition polymer in two ways: its monomers need two functional groups rather than a C=C double bond, and a second product is formed, so its atom economy is below 100 %. The main chain also contains atoms other than carbon, the oxygen of an ester link or the nitrogen of an amide link.

Polyesters

A polyester contains ester links along its chain. The best known is PET (poly(ethylene terephthalate)), made from benzene-1,4-dicarboxylic acid and ethane-1,2-diol. It is strong, transparent and only slightly permeable to carbon dioxide, so it is used for drinks bottles, and drawn into fibres it is the "polyester" of clothing.

For a long chain the amount of water released is very close to $2n$, and it is often written that way. PLA (polylactic acid) is a polyester made from a single monomer, lactic acid, ${{\mathrm{CH}{\vphantom{X}}_{\smash[t]{3}}\mathrm{CH}(\mathrm{OH})\mathrm{COOH}}}$, obtained by fermenting sugars from plants such as maize or sugar cane. Its repeating unit is ${{{-}[\mathrm{O}{-}\mathrm{CH}(\mathrm{CH}{\vphantom{X}}_{\smash[t]{3}}){-}\mathrm{CO}]{\vphantom{X}}_{\smash[t]{n }}{-}}}$. The ester links can be hydrolysed, so PLA cups and packaging break down under industrial composting conditions.

Polyamides

A polyamide contains amide links along its chain. Nylon-6,6, first made by Wallace Carothers at DuPont in 1935, is formed from hexane-1,6-diamine and hexanedioic acid. Hydrogen bonds form between the N–H of one chain and the C=O of the next, which makes nylon strong and gives it a high melting point; it is used for ropes, fishing lines, clothing and engineering parts such as gears.

In the laboratory, nylon can be made at room temperature by carefully layering a solution of hexanedioyl dichloride over an aqueous solution of hexane-1,6-diamine. The polymer forms at the boundary between the two liquids and can be pulled out as a continuous thread, releasing hydrogen chloride.

Kevlar is a polyamide made from benzene-1,4-diamine and benzene-1,4-dicarbonyl dichloride, releasing hydrogen chloride. Its flat, rigid chains line up side by side and are held together by many hydrogen bonds, which makes Kevlar fibres exceptionally strong for their mass. It is used in protective vests, helmets and ropes.

Drawing the repeating unit of a condensation polymer

1. Identify the two functional groups on each monomer and the link they form (ester or amide).

2. Remove the atoms that leave as the small molecule: the OH from each carboxylic acid group (or Cl from each acyl chloride), and an H from each alcohol or amine group.

3. Join the residues of the two monomers with one link, keeping the rest of each monomer unchanged between its two ends.

4. Put the unit containing one residue of each monomer in brackets, with continuation bonds through the brackets and $n$ outside. One end of the unit ends in C=O and the other in the O or N of the next link.

Identifying the monomers of a condensation polymer

To work backwards, find each amide or ester link in the repeating unit and split it, adding back the atoms of the small molecule: OH to the C=O carbon (to make ${{{-}\mathrm{COOH}}}$) and H to the N or O (to make ${{{-}\mathrm{NH}{\vphantom{X}}_{\smash[t]{2}}}}$ or ${{{-}\mathrm{OH}}}$).

Addition and condensation polymers compared

Biological macromolecules: condensation and hydrolysis

The large molecules of living things are all condensation polymers. Proteins are polyamides built from amino acids, polysaccharides such as starch and cellulose are built from sugars such as glucose, and nucleic acids (DNA and RNA) are built from nucleotides. In each case a molecule of water is released as each link forms.

Two amino acids join to form a dipeptide: the carboxyl group of one reacts with the amino group of the other, releasing water and forming an amide link, called a peptide link in proteins. Each end of the dipeptide still carries a functional group, so further amino acids can add on to build a protein chain hundreds of residues long.

Hydrolysis breaks biological macromolecules back into their monomers. During digestion, enzymes catalyse the hydrolysis of proteins to amino acids and of starch to glucose; in the laboratory the same reactions occur when the polymer is heated with aqueous acid. For starch, approximately:

Synthetic polyesters and polyamides can be hydrolysed in the same way, by acids, alkalis or enzymes from microorganisms. This is why condensation polymers can be designed to be biodegradable, while addition polymers, with only C–C and C–H bonds in the chain, cannot.

Applying the ideas in exam questions

Paper 1A: the repeating unit of an addition polymer

Paper 1A: an alloy on the bonding triangle

Paper 1B: identifying a plastic from its density

(Paper 1B style) A student identifies a sample of plastic pellets from its density. The pellets float in water, so a thin wire is used to hold them under the surface while their volume is measured by displacement in a $10\ \text{cm}^3$ measuring cylinder. The student's results are in the table.

Paper 2: two chlorides on the bonding triangle

Paper 2: materials in a bicycle

Paper 2: polyesters and polyamides (HL)

Self-check

• To do:

  I can explain why bonding is better described as a continuum than as three separate types.

• To do:

  I can describe the axes, corners and regions of the bonding triangle.

• To do:

  I can calculate the electronegativity difference and average for a binary compound and place it on the triangle.

• To do:

  I can read the approximate percentage covalent and ionic character of a compound from the triangle.

• To do:

  I can predict and explain the properties of a substance from its position on the triangle, taking its structure into account.

• To do:

  I can describe how the bonding in the period 3 oxides changes across the period.

• To do:

  I can outline the limitations of the bonding triangle and of discrete bonding categories.

• To do:

  I can explain why alloys are mixtures rather than compounds.

• To do:

  I can explain, using non-directional metallic bonding, why alloys form and why they are harder, stronger and less malleable than pure metals.

• To do:

  I can explain why alloys conduct electricity and give examples of alloys with enhanced properties.

• To do:

  I can define polymer, monomer and repeating unit, and give examples of natural and synthetic polymers.

• To do:

  I can explain the common properties of plastics, and the effects of chain length, branching, polar groups and cross-linking, in terms of structure.

• To do:

  I can explain why most plastics are not biodegradable and outline the problems of plastic waste.

• To do:

  I can describe addition polymerization and explain why its atom economy is 100 %.

• To do:

  I can draw the repeating unit of an addition polymer from its monomer, with continuation bonds, and deduce the monomer from a polymer.

• To do:

  I can describe condensation polymerization and identify the small molecule released (HL).

• To do:

  I can draw the repeating units of polyesters and polyamides from given monomers, and deduce the monomers from a repeating unit (HL).

• To do:

  I can classify a polymer as addition or condensation, giving a reason (HL).

• To do:

  I can outline how biological macromolecules form by condensation and break down by hydrolysis (HL).

Open it. Replace C=C with C–C, keep both groups on the second carbon and add a continuation bond to each carbon. The large ester group is a side group; it is not part of the main chain.

As a condensed formula: ${{{-}[\mathrm{CH}{\vphantom{X}}_{\smash[t]{2}}{-}\mathrm{C}(\mathrm{CH}{\vphantom{X}}_{\smash[t]{3}})(\mathrm{COOCH}{\vphantom{X}}_{\smash[t]{3}})]{\vphantom{X}}_{\smash[t]{n }}{-}}}$.

Monomer: restore the double bond between the two carbons, giving ${{\mathrm{CH}{\vphantom{X}}_{\smash[t]{2}}{=}\mathrm{C}(\mathrm{CH}{\vphantom{X}}_{\smash[t]{3}})\mathrm{Cl}}}$, 2-chloropropene.

Join the residues through an amide link, ${{{-}\mathrm{NH}{-}\mathrm{CO}{-}}}$:

The unit contains one residue of each monomer. Both ends carry continuation bonds: the CO at the right end joins the NH of the next unit.

(c)

• PLA contains ester links/polar bonds that can be hydrolysed «by enzymes/microorganisms» ✓

• poly(ethene) has only strong/non-polar C–C and C–H bonds/no functional groups that can be attacked ✓

(d)

• ${{\mathrm{H}{\vphantom{X}}_{\smash[t]{2}}\mathrm{N}{-}(\mathrm{CH}{\vphantom{X}}_{\smash[t]{2}}){\vphantom{X}}_{\smash[t]{6}}{-}\mathrm{NH}{\vphantom{X}}_{\smash[t]{2}}}}$ ✓

• ${{\mathrm{HOOC}{-}(\mathrm{CH}{\vphantom{X}}_{\smash[t]{2}}){\vphantom{X}}_{\smash[t]{8}}{-}\mathrm{COOH}}}$

• Accept the diacyl dichloride ${{\mathrm{ClOC}{-}(\mathrm{CH}{\vphantom{X}}_{\smash[t]{2}}){\vphantom{X}}_{\smash[t]{8}}{-}\mathrm{COCl}}}$

(e)

• PVC: addition AND formed by breaking/opening the C=C of chloroethene with no other product ✓

• PET: condensation AND ester links form between functional groups with loss of water ✓