Learners' AlleyDesigned by Khow Willard

Chemical Bonding and Materials

1

Elements, Compounds and Mixtures

1.1Comparing them

	Element	Compound	Mixture
Made of	one kind of atom	two or more elements chemically combined	two or more substances not chemically combined
Ratio	not applicable	fixed ratio	no fixed ratio
Properties	its own	different from its elements	similar to its substances
Formed by	mostly occurs naturally	a chemical reaction	usually physical mixing
Melting and boiling	fixed points	fixed points	over a range of temperatures
Separated by	cannot be broken down by chemical methods	chemical methods only, e.g. thermal decomposition or electrolysis	physical methods, e.g. filtration or distillation

An element can exist as single atoms (helium), molecules (O₂) or a metal lattice (copper).
Not all molecules are compounds: H₂ is an element and CO₂ is a compound.

2

Why Do Atoms Combine?

2.1The noble gases are stable

The noble gases (Group 18) have a full outermost shell, so they are stable and unreactive. This is a noble gas electronic configuration.

Helium has a duplet, 2 electrons in its only shell. Neon and argon have an octet: 8 outer electrons.
Noble gases are monoatomic. They exist as single atoms.

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Why Do Atoms Combine? (continued)

2.2Three ways to a full outer shell

Other atoms combine and form bonds to reach a noble gas electronic configuration.

Lose electrons

Metals form positive ions.

Gain electrons

Non-metals form negative ions.

Share electrons

Non-metals form covalent bonds.

3

Ionic Bonding

3.1Forming ions

Metal atoms lose electrons

Na → Na⁺ + e⁻
Mg → Mg²⁺ + 2e⁻

Non-metal atoms gain electrons

Cl + e⁻ → Cl⁻
O + 2e⁻ → O²⁻

Na (2.8.1) loses 1 electron to become Na⁺ (2.8). Cl (2.8.7) gains 1 electron to become Cl⁻ (2.8.8). Both ions have a noble gas electronic configuration.
Na could lose 1 electron or gain 7, and losing 1 is far more likely. Na⁺ then has 11 protons and 10 electrons, so its charge is +1.
A negative ion of a single element ends in -ide: chloride, oxide. Metal ions keep the name of the metal.

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3

Ionic Bonding (continued)

3.2Common ions

Charge	Positive ions (cations)	Negative ions (anions)
1	hydrogen H⁺, sodium Na⁺, potassium K⁺, silver Ag⁺, ammonium NH₄⁺	fluoride F⁻, chloride Cl⁻, bromide Br⁻, iodide I⁻, hydroxide OH⁻, nitrate NO₃⁻
2	magnesium Mg²⁺, calcium Ca²⁺, barium Ba²⁺, iron(II) Fe²⁺, copper(II) Cu²⁺, zinc Zn²⁺	oxide O²⁻, carbonate CO₃²⁻, sulfate SO₄²⁻
3	aluminium Al³⁺, iron(III) Fe³⁺	phosphate PO₄³⁻

NoteIons with more than one atom, such as NH₄⁺ and SO₄²⁻, are polyatomic ions. Learn the names, formulae and charges of the five marked like this by heart.

The negative ions of Group 17 are halide ions. Ions with a charge above +3 are rare.
Hydrogen is a non-metal that forms a positive ion. H⁺ has no electrons: it is just a proton.

3.3The ionic bond

An ionic bond is the mutual electrostatic attraction between ions of opposite charges. It usually forms between a metal and a non-metal.

sodium chloride, NaCl

Show only the outer electrons, as dots for one atom and crosses for the other.
Put each ion in square brackets, with its charge outside.
An ionic compound is neutral: total positive charge = total negative charge.

Worked exampleMagnesium reacts with iodine. Give the formula of the compound and draw its dot-and-cross diagram.

Magnesium forms Mg²⁺ (charge +2). Iodine forms I⁻ (charge −1).To balance, one Mg²⁺ needs two I⁻, as +2 + 2(−1) = 0.The formula is MgI₂. Show the number of each ion in front of its bracket.

Exam tipAn ionic formula shows the simplest ratio of ions, with no charges.

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Ionic Bonding (continued)

3.4The giant ionic lattice

In the solid state, an ionic compound is a giant ionic lattice: a three-dimensional structure of alternating positive and negative ions in a regular, repeating pattern.

Every ion attracts all the oppositely charged ions next to it. In sodium chloride, each Na⁺ is surrounded by six Cl⁻, and each Cl⁻ by six Na⁺.

3.5Properties of ionic compounds

Property	Explanation
High melting and boiling points, solid at room temperature	The electrostatic attraction between oppositely charged ions is strong, so a large amount of energy is needed to overcome it.
Hard but brittle	A strong force moves the layers so that ions of the same charge meet. They repel and the lattice shatters.
Usually soluble in water, insoluble in organic solvents	Water pulls the ions out of the lattice.
Conduct electricity when molten or aqueous, not as a solid	The ions are mobile only when molten or dissolved. In the solid they only vibrate about fixed positions.

Worked exampleExplain why sodium chloride has a high melting point.

Sodium chloride has a giant ionic lattice structure.There are strong electrostatic forces of attraction between the oppositely charged Na⁺ and Cl⁻ ions.A large amount of energy is needed to overcome these forces, so the melting point is high.

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4

Covalent Bonding

4.1Sharing electrons

DefinitionA covalent bond is formed when two atoms share a pair of electrons (a bonding pair), so each atom gains a noble gas electronic configuration. It forms between non-metal atoms.

Valency is the number of electrons an atom must lose, gain or share to reach a noble gas electronic configuration. For a non-metal, the bonds it forms = 8 − number of valence electrons (hydrogen: 2 − 1 = 1).

Element	Group	Valence electrons	Valency
hydrogen	1	1	1
carbon, silicon	14	4	4
nitrogen, phosphorus	15	5	3
oxygen, sulfur	16	6	2
fluorine, chlorine, bromine, iodine	17	7	1

4.2Dot-and-cross diagrams of molecules

Cl₂: single bond, Cl–Cl

O₂: double bond, O=O

N₂: triple bond, N≡N

water, H₂O

ammonia, NH₃

methane, CH₄

carbon dioxide, CO₂

Draw each shared pair between the two symbols, and the other outer electrons in pairs.
Cl₂, O₂ and N₂ are diatomic: each molecule has two atoms. A structural formula shows each shared pair as a line, as in Cl–Cl, O=O and N≡N.

Worked exampleDraw a dot-and-cross diagram for hydrogen sulfide, H₂S.

Hydrogen has valency 1. Sulfur has 6 outer electrons, so its valency is 2.Sulfur forms two single bonds, one to each hydrogen atom.Sulfur keeps its other 4 outer electrons as 2 lone pairs.

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Covalent Bonding (continued)

4.3Properties of simple covalent substances

Strong covalent bonds inside each molecule.
Weak intermolecular forces of attraction between the molecules.
Larger molecules have stronger forces between them, so higher melting and boiling points. At room temperature methane is a gas, bromine a liquid and iodine a solid.
At low temperatures, some simple molecules form a lattice, as water does in ice.

Property	Explanation
Low melting and boiling points, often volatile	A small amount of energy overcomes the weak intermolecular forces.
Usually insoluble in water but soluble in organic solvents	They are neutral molecules.
Do not conduct electricity	No mobile ions or electrons. Hydrogen chloride conducts only after it dissolves in water and forms H⁺ and Cl⁻.

Exam tipWhen a simple covalent substance melts or boils, the weak intermolecular forces between molecules are overcome. The covalent bonds inside the molecules do not break.

5

Giant Covalent Structures

5.1Simple and giant molecules

Simple molecules

A countable number of atoms in a fixed ratio, e.g. H₂O, CO₂ and Br₂.

Giant molecules

An uncountable number of atoms all joined by covalent bonds, e.g. diamond and silicon dioxide.

5.2Diamond, graphite and silicon dioxide

Allotropes are different forms of the same element with different arrangements of atoms. Diamond and graphite are allotropes of carbon.

diamond

graphite

silicon dioxide: Si larger, O smaller

	Diamond	Graphite	Silicon dioxide
Bonding	each C bonded to 4 C, tetrahedral	each C bonded to 3 C, in layers of hexagons	each Si bonded to 4 O, SiO₂, as in quartz and sand
Hardness	very hard	soft and slippery	hard
Conducts?	no	yes	no

Diamond coats drill bits and cutting tools. Graphite is used as a solid lubricant and in pencil lead.

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5

Giant Covalent Structures (continued)

5.3Properties of giant covalent substances

Property	Explanation
Very high melting and boiling points, solids	Many strong covalent bonds link all the atoms, so a large amount of energy is needed to break them.
Graphite is soft and slippery	Only weak forces hold the layers together, so the layers slide over each other.
Graphite conducts electricity	Each carbon uses 3 of its 4 outer electrons in bonds. The fourth is delocalised and moves along the layers.
Diamond and silicon dioxide do not conduct	All outer electrons are used in bonds, so there are no mobile electrons.

Giant covalent substances are insoluble in water and in organic solvents.

NoteGraphene (a single layer of graphite), carbon nanotubes and fullerenes are other allotropes of carbon.

5.4Macromolecules

A polymer is many small molecules joined into long chains, e.g. poly(ethene), nylon, starch and silk.
The chains have different lengths, so polymers have no fixed melting point and soften gradually on heating.
Most do not conduct electricity. Most are insoluble in water but soluble in organic solvents.

6

Metallic Bonding and Alloys

6.1A sea of delocalised electrons

A metallic bond is the mutual electrostatic attraction between the positive metal ions and the sea of delocalised electrons.

Metal atoms lose their outer electrons and form a giant metallic lattice of positive ions.
The electrons no longer belong to any one atom. They are delocalised and move freely between the ions.
Metal atoms form only positive ions, so ionic bonding between them is unlikely.

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Metallic Bonding and Alloys (continued)

6.2Properties of metals

Property	Explanation
High melting and boiling points	The strong attraction between the ions and the delocalised electrons needs a large amount of energy to overcome.
Conduct electricity as a solid and when molten	The delocalised electrons are mobile and carry the current.
Conduct heat	The delocalised electrons pass on thermal energy through the lattice.
Malleable and ductile	The layers of identical ions slide over one another, and the sea of electrons keeps holding them together.

6.3Alloys

An alloy is a mixture of a metal with one or more other elements.

Different-sized atoms disrupt the orderly layers, so the layers cannot slide easily. Alloys are harder and stronger than the pure metal.
An alloy is a mixture, so it melts over a range of temperatures. Its make-up is given as approximate percentages by mass, not as a formula.

Worked exampleBrass is an alloy of copper and zinc. Explain why brass is harder than pure copper.

In pure copper, the layers of identical copper ions can slide over one another.In brass, the zinc atoms are a different size and disrupt the orderly layers.The layers cannot slide over one another easily, so brass is harder.

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6

Metallic Bonding and Alloys (continued)

6.4Uses of alloys

Alloy	Made from	Why it is used
bronze	copper and tin	harder and stronger, e.g. statues and bells
brass	copper and zinc	harder than copper
solder	tin and lead	lower melting point
steel	iron and carbon	harder than iron
stainless steel	iron, carbon, chromium, nickel	resists corrosion

7

Structure and Properties

7.1The three types of bond

	Ionic	Covalent	Metallic
Between	metal and non-metal	non-metals	metal atoms
Electrons	transferred	shared	delocalised
Held by attraction between	oppositely charged ions	a shared pair of electrons and both nuclei	positive ions and delocalised electrons

Exam tipTo state the type of bonding, look at the elements: a metal and a non-metal are ionic, non-metals only are covalent, and a metal element is metallic.

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Structure and Properties (continued)

7.2Which type of structure is it?

	Ionic	Simple covalent	Giant covalent	Metallic
Particles	ions	molecules	atoms	ions and electrons
Melting point	high	low	very high	high
Conducts	molten, aqueous	no	no (graphite yes)	solid and molten
Soluble in water	usually	usually not	no	no

1. Is the melting point low? Yes: a simple covalent substance. No: go on.

2. Does it conduct as a solid? Yes: a metal or graphite. No: go on.

3. Does it conduct when molten or aqueous? Yes: an ionic compound. No: a giant covalent substance.

NoteA substance that melts over a range of temperatures is a mixture, such as an alloy, or a polymer.

Worked exampleSolid A is an insulator but conducts when dissolved in water. Solid B is insoluble in water but conducts as a solid. What bonding do they have?

A conducts only when dissolved, so its ions are mobile in solution. A has ionic bonding.B conducts as a solid, so it has mobile electrons. B has metallic bonding.

Worked exampleSolids X, Y and Z are graphite, sodium chloride and silicon dioxide. X is soft and slippery and conducts electricity. Y and Z have very high melting points. Y conducts when dissolved in water but not as a solid. Z does not conduct. Identify X, Y and Z.

X conducts as a solid and is slippery, so X is graphite.Y conducts only when dissolved, so it is ionic: Y is sodium chloride.Z has a very high melting point but never conducts, so it is giant covalent: Z is silicon dioxide.

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