Chemical Bonding

Structure and Properties
of Materials

IP Chemistry (6092), Sec 3

By the end of this topic, you can answer

1

Why do some atoms lose electrons, some gain them and some share them?

2

How do sodium and chlorine, a dangerous metal and a poisonous gas, make safe table salt?

3

How are the electrons arranged in a water molecule, and why is its formula H₂O?

4

Why does salt shatter when hit, but a metal spoon only bends?

5

Diamond and pencil lead are both carbon. Why is one the hardest substance and the other soft?

6

How can you tell what kind of substance something is from its properties?

Starter

PredictDiamond is the hardest natural substance. The lead in your pencil is soft enough to rub off on paper. Both are made of only carbon atoms.

How can two substances made of the same atoms be so different?

NoteThe answer is in how the atoms are joined. This topic covers the three types of bond, and how the bonding and structure of a substance decide its properties.
1

Elements, Compounds and Mixtures

1.1Comparing them

ElementCompoundMixture
Made ofone kind of atomtwo or more elements chemically combinedtwo or more substances not chemically combined
Rationot applicablefixed rationo fixed ratio
Formed byoccurs naturallya chemical reactionusually physical mixing
Melting and boilingfixed pointsfixed pointsover a range of temperatures
1

Elements, Compounds and Mixtures

1.2Telling them apart

An element cannot be broken down into simpler substances by chemical methods. It can be atoms (helium), molecules (O2) or a metal lattice (copper).
  • Not all molecules are compounds: H2 is an element and CO2 is a compound.
  • A compound splits into elements by thermal decomposition or electrolysis.
CompoundA water molecule near the top of a glass is identical to one at the bottom. The ratio of hydrogen to oxygen is always 2 : 1.
MixtureAir on a mountain top has less oxygen than air at sea level. Mixtures are separated by physical methods.
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.
2

Why Do Atoms Combine?

2.2Three ways to a full outer shell

Most atoms do not have a full outermost shell. To get one, atoms chemically combine and form bonds. They can do this in three ways.

Lose electronsForms a positive ion. Metals do this.
Gain electronsForms a negative ion. Non-metals do this.
Share electronsForms a covalent bond between non-metal atoms.
NoteHydrogen has only 1 electron in its only shell, so it reacts to fill the shell. This is why hydrogen balloons are dangerous but helium balloons are not.
3

Ionic Bonding

3.1Forming ions

Metal atoms lose electrons. Non-metal atoms gain electrons. The ions have a noble gas configuration.

3

Ionic Bonding

3.1Cations and anions

Positive ion (cation)An atom loses electrons. It has more protons than electrons, so it is positively charged.
Negative ion (anion)An atom gains electrons. It has more electrons than protons, so it is negatively charged.
  • A sodium atom (2.8.1) loses 1 electron to become Na+ (2.8). Losing 1 is far more likely than gaining 7.
  • A chlorine atom (2.8.7) gains 1 electron to become Cl− (2.8.8).
3

Ionic Bonding

3.1Writing ion formation

Na → Na+ + e−
Mg → Mg2+ + 2e−
Cl + e− → Cl−
O + 2e− → O2−
  • A negative ion of a single element ends in -ide: chloride, oxide. Metal ions keep the name of the metal.
  • Ions with a charge above +3 are rare, because giving up more electrons becomes too difficult.
3

Ionic Bonding

3.2Common ions

ChargePositive ions (cations)Negative ions (anions)
1hydrogen H+, sodium Na+, potassium K+, silver Ag+, ammonium NH4+fluoride F−, chloride Cl−, bromide Br−, iodide I−, hydroxide OH−, nitrate NO3−
2magnesium Mg2+, calcium Ca2+, barium Ba2+, iron(II) Fe2+, copper(II) Cu2+, zinc Zn2+oxide O2−, carbonate CO32−, sulfate SO42−
3aluminium Al3+, iron(III) Fe3+phosphate PO43−
NoteIons with more than one atom are polyatomic ions. The negative ions of Group 17 are halide ions. The five marked like this are used throughout the O Level syllabus, so learn their names, formulae and charges by heart.
3

Ionic Bonding

sodium chloride, NaCl

3.3The ionic bond

An ionic bond is the mutual electrostatic attraction between ions of opposite charges.
  • Show only the outer electrons, as dots for one atom and crosses for the other.
  • Put each ion in square brackets, charge outside.
  • Ionic bonds usually form between a metal and a non-metal.
AnalogyOpposite charges pull like the north and south poles of two magnets. Each ion is held by that pull, not by any shared electrons.
3

Ionic Bonding

3.3Ionic compounds and their formulae

An ionic compound is a neutral substance made of ions held together by ionic bonds. The total positive charge equals the total negative charge.

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

Step 1Magnesium forms Mg2+. Iodine forms I−.

Step 2One Mg2+ needs two I−, as +2 + 2(−1) = 0.

Step 3The formula is MgI2

3

Ionic Bonding

3.3Working out a formula

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

Step 1Sodium forms Na+ and oxygen forms O2−.

Step 2Two Na+ are needed for one O2−.

Step 3The formula is Na2O

Exam tipAn ionic formula shows the simplest ratio of ions, with no charges. Check that the total positive charge equals the total negative charge.
3

Ionic Bonding

Quick checkAluminium has 3 electrons in its outer shell and oxygen forms O²⁻. What is the formula of aluminium oxide?

AAlO
BAlO₂
CAl₂O₃
DAl₃O₂
NoteC: aluminium forms Al³⁺ and oxygen forms O²⁻. Two Al³⁺ give +6 and three O²⁻ give −6, so the formula is Al₂O₃.
3

Ionic Bonding

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

Ionic Bonding

3.5Properties of ionic compounds

PropertyWhy (giant ionic lattice)
High melting and boiling pointsThe electrostatic attraction between oppositely charged ions is strong, so a large amount of energy is needed to overcome it.
Hard but brittleA strong force pushes layers so that like charges meet. They repel and the lattice shatters.
Usually soluble in waterWater pulls the ions out of the lattice. Organic solvents cannot.
Conduct when molten or aqueousThe ions are mobile only when molten or dissolved. In the solid they vibrate in fixed positions.
3

Ionic Bonding

3.5Why ionic compounds are brittle

PredictSalt is hard, yet it shatters when hit. What happens when a layer of ions is pushed along by one ion?

When enough force is applied, ions of the same charge approach each other. The repulsion is larger than the attraction, so the lattice breaks.

3

Ionic Bonding

3.5Electrical conductivity of sodium chloride

PredictSolid salt does not conduct electricity, but molten salt does. What must be different?

SolidIons fixed in place, so no current.
MoltenIons move freely, so it conducts.
AqueousIons move freely in the water, so it conducts.
3

Ionic Bonding

Worked exampleExplain why sodium chloride has a high melting point.

Step 1Sodium chloride has a giant ionic lattice structure.

Step 2There are strong electrostatic forces of attraction between the oppositely charged Na+ and Cl− ions.

Step 3A large amount of energy is needed to overcome these forces, so the melting point is high.

3

Ionic Bonding

Quick checkSolid X does not conduct electricity but conducts when dissolved in water. What type of substance is X most likely to be?

Aa metal
Ban ionic compound
Cgraphite
Da giant covalent compound
NoteB: the ions cannot move in the solid, but are free to move in solution. A metal and graphite conduct as solids. A giant covalent compound does not conduct at all.
4

Covalent Bonding

4.1Sharing electrons

DefinitionA covalent bond is formed when two atoms share a pair of electrons (a bonding pair). It gives each atom a noble gas electronic configuration.
  • It forms between non-metal atoms. Both nuclei attract the shared pair, so the bond is hard to break.
  • Draw each shared pair between the two symbols, and the other outer electrons in pairs.
AnalogyA shared pair is like a rope that two people both grip.
4

Covalent Bonding

4.1How many bonds does an atom form?

Valency is the number of electrons an atom must lose, gain or share to reach a noble gas electronic configuration. For a non-metal it is 8 − the number of valence electrons (hydrogen: 2 − 1 = 1).
ElementGroupValence electronsValency (electrons shared)
hydrogen112 − 1 = 1
carbon, silicon1448 − 4 = 4
nitrogen, phosphorus1558 − 5 = 3
oxygen, sulfur1668 − 6 = 2
fluorine, chlorine, bromine, iodine1778 − 7 = 1
4

Covalent Bonding

4.2Single, double and triple bonds

Cl2: single bondeach atom shares 1 electron
O2: double bondeach atom shares 2 electrons
N2: triple bondeach atom shares 3 electrons
  • These molecules are diatomic: each has two atoms. A structural formula shows each shared pair as a line, as in Cl–Cl, O=O and N≡N.
4

Covalent Bonding

4.2Molecules of compounds

water, H2O
ammonia, NH3
methane, CH4
carbon dioxide, CO2
  • Match the sharing: hydrogen shares 1 electron and oxygen shares 2, so one oxygen atom bonds to two hydrogen atoms in water.
  • In CO2, carbon forms a double bond with each oxygen atom.
4

Covalent Bonding

4.2Drawing a dot-and-cross diagram

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

Step 1Hydrogen has 1 outer electron and shares 1. Sulfur has 6 and shares 2.

Step 2So sulfur forms two single bonds, one to each hydrogen atom.

Step 3Sulfur keeps its other 4 outer electrons as 2 lone pairs.

4

Covalent Bonding

Quick checkNitrogen forms molecules with the formula N₂. How many electrons does each nitrogen atom share?

A1
B2
C3
D5
NoteC: nitrogen has 5 outer electrons and valency 8 − 5 = 3. Two nitrogen atoms share three pairs and form a triple bond, N≡N.
4

Covalent Bonding

4.3Properties of simple covalent substances

PredictBromine boils at only 59 °C. When it boils, do its molecules fall apart?

  • Strong covalent bonds inside each molecule.
  • Weak intermolecular forces between the molecules.
  • Melting only separates the molecules. It does not break the covalent bonds.
AnalogyEach molecule is a pair of clasped hands. The molecules are easy to pull apart, but the hands stay clasped.
4

Covalent Bonding

4.3Properties of simple covalent substances

PropertyExplanation
Low melting and boiling points, often volatileA small amount of energy overcomes the weak intermolecular forces. Larger molecules have stronger forces, so higher melting points.
Usually insoluble in water but soluble in organic solventsThey are neutral molecules.
Do not conduct electricityNo mobile ions or electrons. Hydrogen chloride conducts only after it dissolves in water and forms H+ and Cl−.
4

Covalent Bonding

4.3Size of molecule and state

SubstanceState at r.t.p.
hydrogen, H2gas
methane, CH4gas
pentane, C5H12liquid
bromine, Br2liquid
iodine, I2solid
NoteLarger molecules have stronger forces between them, so their melting and boiling points are higher.
4

Covalent Bonding

Quick checkMothballs are made of naphthalene, a covalent molecule. It melts at 80 °C. Potassium chloride melts at 770 °C. Which statement is correct?

ANaphthalene has strong ionic bonds between its molecules.
BNaphthalene has weak forces between its molecules, so it melts easily.
CPotassium chloride has weak forces between its ions.
DBoth have giant structures.
NoteB: only a little energy overcomes the weak forces between the molecules. Potassium chloride is a giant ionic lattice, which needs a lot of energy to melt.
5

Giant Covalent Structures

5.1Simple molecules and giant molecules

Simple moleculesA countable number of atoms in a fixed ratio, e.g. H2O, CO2 and wax, C18H38.
Giant moleculesAn uncountable number of atoms all joined by covalent bonds, e.g. diamond and sand.
5

Giant Covalent Structures

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.
diamondeach C bonded to 4 C, tetrahedral
graphiteeach C bonded to 3 C, in layers of hexagons
silicon dioxideSi (larger) and O (smaller), SiO2
5

Giant Covalent Structures

5.3Properties of giant covalent substances

Many strong covalent bonds link all the atoms. A large amount of energy breaks them, so melting points are very high.
DiamondGraphite
Hardnessvery hardsoft: only weak forces between its layers, so they slide over each other. Used as a lubricant and in pencil lead.
Conducts electricity?no: all 4 outer electrons are in bondsyes: each carbon uses only 3, leaving one free electron per atom to carry current
Soluble?Giant covalent substances are insoluble in water and organic solvents.
AnalogyGraphite is like a stack of paper sheets. Each sheet is strong, but the sheets slide easily off one another, which is why pencil lead rubs off onto paper.
5

Giant Covalent Structures

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 electrons

A metallic bond is the mutual electrostatic attraction between the positive metal ions and the sea of delocalised electrons.
  • 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.
6

Metallic Bonding and Alloys

6.2Properties of metals

  • 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: the sea of delocalised electrons is mobile in solid and liquid.
  • Conduct heat: the delocalised electrons pass on thermal energy through the lattice.
  • Malleable and ductile (pure metals): layers slide, while the electron sea keeps holding the ions together.
6

Metallic Bonding and Alloys

6.3Alloys

An alloy is a mixture of a metal with one or more other elements.
  • Pure metal: layers of same-sized ions slide over one another, so it is malleable and ductile.
  • Alloy: different-sized atoms disrupt the layers, so alloys are harder and stronger.
6

Metallic Bonding and Alloys

6.4Uses of alloys

AlloyMade fromWhy it is used
bronzecopper and tinharder and stronger, e.g. statues and bells
brasscopper and zincharder than copper
soldertin and leadlower melting point
steeliron and carbonharder than iron
stainless steeliron, carbon, chromium, nickelresists corrosion
NoteThe formula of an alloy is given as percentages by mass, because the exact numbers of atoms are not fixed.
7

Structure and Properties

7.1The three types of bond

IonicCovalentMetallic
Betweenmetal and non-metalnon-metalsmetal atoms
Electronstransferredshareddelocalised
Held by attraction betweenoppositely charged ionsa shared pair of electrons and both nucleipositive ions and delocalised electrons
ExampleNaCl, MgOH2O, CH4Na, Cu
Exam tipName the particles in the bond: ions of opposite charge, a shared pair of electrons, or positive ions and delocalised electrons.
7

Structure and Properties

Worked exampleState the type of bonding in potassium bromide, ammonia and magnesium.

Step 1Potassium is a metal and bromine is a non-metal, so potassium bromide has ionic bonding.

Step 2Nitrogen and hydrogen are both non-metals, so ammonia has covalent bonding.

Step 3Magnesium is a metal element, so it has metallic bonding.

7

Structure and Properties

7.2Which type of substance is it?

7

Structure and Properties

7.2Summary

IonicSimple covalentGiant covalentMetallic
Particlesionsmoleculesatomsions and electrons
Melting pointhighlowvery highhigh
Conductsmolten, aqueousnono (graphite yes)solid and molten
Soluble in waterusuallyusually notnono
7

Structure and Properties

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

Step 1X conducts as a solid and is slippery, so X is graphite.

Step 2Y conducts only when dissolved, so it is ionic: Y is sodium chloride.

Step 3Z has a very high melting point but never conducts, so Z is silicon dioxide.

Exit ticket

1Give the formula of calcium chloride.

CaCl₂[1]

2How many bonds does a carbon atom form in methane?

Four, as carbon has 4 valence electrons and valency 8 − 4 = 4.[1]

3Why can graphite conduct electricity but diamond cannot?

Each carbon in graphite has one free electron that can move. Diamond uses all four outer electrons in bonds.[1]

4Why is an alloy harder than the pure metal?

Atoms of different sizes disrupt the layers, so they cannot slide easily.[1]