Ionic Bonding Notes

Why Atoms Form Bonds?

• Bonds involve electrons in the outer shells of atoms.
• Filled electron shells are very stable.
  • 1st shell holds a maximum of 2 electrons.
  • 2nd shell holds a maximum of 8 electrons.
  • 3rd shell holds a maximum of 8 electrons.
• Each shell has a maximum number of electrons that it can hold.
• Electrons fill the shells nearest the nucleus first.

Noble Gases

• Noble gases have completely full outer shells and are stable (inert).
• Atoms of other elements have incomplete outer electron shells and are unstable.
• Noble gases are very unreactive and do not usually form bonds.
• By forming bonds, atoms of other elements can achieve filled outer shells and become stable.

Types of Bonding

• Different types of bonds are formed depending on the types of atoms involved.
• All bonds involve electrons and all bonding involve changes to the number of electrons in the outer shells of atoms.
  • Ionic bonding (electrovalent): occurs between metal and non-metal atoms.
  • Covalent bonding: occurs between non-metal atoms only.
  • Metallic bonding: occurs between metal atoms only.

Chemical Bonds (Strong)

• Ionic (or electrovalent)
• Covalent
  • Molecular
  • Network (or co-ordinate)
• Metallic

Physical Bonds (Weak)

• Induced dipole-dipole interactions (London forces)
• Permanent dipole-dipole interactions (Term 3) (van der Waals‘ forces)
• Hydrogen bonds (strongest among weak bonds)

From Atoms to Ions

• Reactive metal atoms can become stable positive ions.

How Atoms Form Ions

• An ion is an atom or group of atoms with an electrical charge (positive or negative).
• Atoms have an equal number of protons and electrons and so do not have an overall charge.
• Atoms with incomplete outer electron shells are unstable.
• By either gaining or losing electrons, atoms can obtain full outer electron shells and become stable.
• When this happens, atoms have an unequal number of protons and electrons and so have an overall charge.
• This is how atoms become ions.

Positive Ions (Cations)

• Also known as cations; they are smaller than the original atom.
• Formed when electrons are removed from atoms.
• The energy associated with the process is known as the ionization energy.

1st Ionization Energy (1st I.E.)

• The energy required to remove one mole of electrons (to infinity) from one mole of gaseous atoms to form one mole of gaseous positive ions. E.g.,
  • Na(g) \longrightarrow Na^+(g) + e^-
  • Mg(g) \longrightarrow Mg^+(g) + e^-
• Successive I.E.’s get larger as the proton:electron ratio increases.
• Large jumps in value occur when electrons are removed from shells nearer the nucleus because there is less shielding and more energy is required to overcome the attraction.
• If the I.E. values are very high, covalent bonding will be favored (e.g., beryllium).

Negative Ions (Anions)

• Known as anions.
• Are larger than the original atom due to electron repulsion in the outer shell.
• Formed when electrons are added to atoms.
• Energy is released as the nucleus pulls in an electron.
• This energy is the electron affinity.

Electron Affinity

• The energy change when one mole of gaseous atoms acquires one mole of electrons (from infinity) to form one mole of gaseous negative ions. E.g.,
  • Cl(g) + e^- \longrightarrow Cl^-(g)
  • O(g) + e^- \longrightarrow O^-(g)
• The greater the effective nuclear charge (E.N.C.), the easier an electron is pulled in.

Positive and Negative Ions

• The electron configuration of an atom shows how many electrons it must lose or gain to have a filled outer shell.
• An atom that loses electrons has more protons than electrons and so has a positive overall charge (positive ion).
• An atom that gains electrons has more electrons than protons and so has a negative overall charge (negative ion).
• Atoms with a nearly empty outer shell will lose electrons to obtain a full outer shell.
• Atoms with a nearly full outer shell will gain electrons to obtain a full outer shell.

Formation of Positive Ions

• An atom that loses one or more electrons forms a positive ion.
• Positive ions have a small ‘+’ symbol and a number by this to indicate how many electrons have been lost.
• This number is usually the same as the number of electrons in the atom’s outer shell. For example:
  • Lithium atom (2.1) → Lithium ion (2) = Li^+
  • Magnesium atom (2.8.2) → Magnesium ion (2.8) = Mg^{2+}
  • Aluminum atom (2.8.3) → Aluminum ion (2.8) = Al^{3+}

Sodium Ion Formation

• Sodium atom: 2.8.1 (partially full outer shell)
  • 11 protons = +11
  • 11 electrons = -11
  • Total charge = 0
• Sodium ion: [2.8] (full outer shell)
  • 11 protons = +11
  • 10 electrons = -10
  • Total charge = +1
• Na \longrightarrow Na^+ (loses 1 electron)

Magnesium Ion Formation

• Magnesium atom: 2.8.2 (partially full outer shell)
  • 12 protons = +12
  • 12 electrons = -12
  • Total charge = 0
• Magnesium ion: [2.8]^{2+} (full outer shell)
  • 12 protons = +12
  • 10 electrons = -10
  • Total charge = +2
• Mg \longrightarrow Mg^{2+} (loses 2 electrons)

Sulfide Ion Formation

• Sulfur atom: 2.8.6 (partially full outer shell)
  • 16 protons = +16
  • 16 electrons = -16
  • Total charge = 0
• Sulfide ion: [2.8.8]^{2-} (full outer shell)
  • 16 protons = +16
  • 18 electrons = -18
  • Total charge = -2
• S \longrightarrow S^{2-} (gains 2 electrons)

Ionic Bonding

• Sodium chloride is an ionic compound formed by the reaction between the metal sodium and the non-metal chlorine.
  • Chlorine has 7 electrons in its outer shell (2.8.7).
• By gaining an electron from sodium, it has a filled outer shell and forms a negative ion.
• Cl + e^- \longrightarrow Cl^- [2.8.7] [2.8.8]
• Na \longrightarrow Na^+ + e^- 2.8.1 [2.8]
• The resulting ions are held together in a crystal lattice by electrostatic attraction.

Formation of Ionic Bonds

• The positive sodium ions and the negative chloride ions are strongly attracted to each other.
• This electrostatic attraction forms ionic bonds in sodium chloride and other ionic compounds.

Formation of Magnesium Chloride

• Mg \longrightarrow Mg^{2+} + 2e^-
• 2Cl + 2e^- \longrightarrow 2Cl^-

Formula of Aluminum Bromide

• What is the formula of aluminum bromide? AlBr3
  • Al^{+3}
  • Br^{-1}
  • 3 bromide ions are needed for each aluminium ion (1:3).

Formula of Aluminum Oxide

• What is the formula of aluminium oxide? Al2O3
  • Al^{+3}
  • O^{-2}
  • 2 aluminium ions are needed for 3 oxide ions (2:3).

Sodium Chloride Formation

• Na \longrightarrow Na^+ + e^-
  • 2,8,1 2,8
• Cl + e^- \longrightarrow Cl^-
  • 2,8,7 2,8,8
• Electron Transferred

Giant Ionic Crystal Lattice

• Oppositely charged ions held in a regular 3-dimensional lattice by electrostatic attraction.
• The arrangement of ions in a crystal lattice depends on the relative sizes of the ions.
• The Na^+ ion is small enough relative to a Cl^- ion to fit in the spaces so that both ions occur in every plane.

Heating Ionic Compounds

• Ionic compounds are solid at room temperature and have high melting points and boiling points because strong ionic bonds hold ions together.
• Larger ionic charges produce stronger ionic bonds, so more heat is required to break the ionic bonds in magnesium oxide than in sodium chloride.
  • Sodium chloride: 1+ and 1- charge; Melting point 801°C and Boiling point 1,413°C
  • Magnesium oxide: 2+ and 2- charge; Melting point 2,852°C and Boiling point 3,600°C

Formation of Ionic Compound Crystals

• All ionic compounds form lattices and crystals when solid.

Electrical Conductivity of Ionic Compounds

• As solids, ionic compounds cannot conduct electricity because their ions are bonded together in the lattice.
• When liquid (molten), the ions can break free of the lattice and are able to move. The ions are charged particles and so can carry an electric current.

Electrical Properties of Ionic Compounds

• Solid ionic compounds do not conduct electricity because ions are held strongly together and cannot move to the cathode or anode.
• Molten ionic compounds do conduct electricity because ions have more freedom in a liquid and can move.
• Solutions of ionic compounds in water do conduct electricity because dissolving an ionic compound in water breaks up the structure, so ions are free to move to the electrodes.

Exam Question

• Explain, in terms of structure and bonding, why strontium chloride can conduct electricity when molten. Use a relevant chemical equation to support your answer. (3 marks)
• Marking Key:
  • When molten, the ions in SrCl_2 dissociate. (1 mark)
  • SrCl_2(s) \longrightarrow Sr^{2+}(l) + 2Cl^-(l) (1 mark)
  • This produces mobile charge in the form of freely moving ions, thus allowing the liquid to conduct electricity. (1 mark)

Brittleness of Ionic Compounds

• Ionic compounds are brittle and shatter when hit.
• When the lattice is hit, a layer of ions is shifted so that ions with the same charges are lined up together.
• These like charges repel each other and so split the ionic lattice, causing it to shatter.
• Ionic bonds are strong, so why does this happen?

Ionic Lattices

• If you move a layer of ions with a force, you get ions of the same charge next to each other.
• The layers repel each other, and the crystal breaks up.

Physical Properties of Ionic Compounds

• Melting point: Very high because a large amount of energy must be put in to overcome the strong electrostatic attractions and separate the ions.
• Strength: Very brittle because any dislocation leads to the layers moving and similar ions being adjacent. The repulsion splits the crystal.
• Electrical conductivity: Do not conduct when solid because ions are held strongly in the lattice; conduct when molten or in aqueous solution because the ions become mobile and conduction takes place.
• Solubility: Insoluble in non-polar solvents but soluble in water. Water is a polar solvent and stabilizes the separated ions. Much energy is needed to overcome the electrostatic attraction and separate the ions, but stability attained by being surrounded by polar water molecules compensates for this.

Ionic Bonds

• Non-metals are highly electronegative (tendency to attract electrons), and metals have low ionization energy; therefore, the valence electrons get donated to the non-metals forming ions.
• Now they form cations and anions - electro static attraction is ionic bonding.

Metallic Bonding

• Involves a lattice of positive ions surrounded by delocalized electrons.
• Metal atoms achieve stability by “off-loading” electrons to attain the electronic structure of the nearest noble gas.
• These electrons join up to form a mobile cloud which prevents the newly-formed positive ions from flying apart due to repulsion between similar charges.
• Atoms arrange in regular close packed 3-dimensional crystal lattices.
• The outer shell electrons of each atom leave to join a mobile “cloud” or “sea” of electrons which can roam throughout the metal.
• The electron cloud binds the newly-formed positive ions together.

Metallic Bond Strength

• Depends on the number of outer electrons donated to the cloud and the size of the metal atom/ion.
• The strength of the metallic bonding in sodium is relatively weak because each atom donates one electron to the cloud.
• The metallic bonding in potassium is weaker than in sodium because the resulting ion is larger and the electron cloud has a bigger volume to cover so is less effective at holding the ions together.
• The metallic bonding in magnesium is stronger than in sodium because each atom has donated two electrons to the cloud. The greater the electron density holds the ions together more strongly.
• Na < Mg > K

Metallic Properties

• Mobile electron cloud allows the conduction of electricity.
• For a substance to conduct electricity, it must have mobile ions or electrons.
• Because the ELECTRON CLOUD IS MOBILE, electrons are free to move throughout its structure.
• Electrons attracted to the positive end are replaced by those entering from the negative end.
• Metals are excellent conductors of electricity.

Malleability and Ductility

• Malleable - Can be hammered into sheets (compressive stress).
• Ductile - Can be drawn into rods and wires (tensile stress).
• As the metal is beaten into another shape, the delocalized electron cloud continues to bind the “ions” together.
• Some metals, such as gold, can be hammered into sheets thin enough to be translucent.
• Metals can have their shapes changed relatively easily.

High Melting Points

• Melting point is a measure of how easy it is to separate individual particles.
• In metals, it is a measure of how strong the electron cloud holds the + ions.
• The ease of separation of ions depends on the:
  • ELECTRON DENSITY OF THE CLOUD
  • IONIC / ATOMIC SIZE

Periods

• Na (2,8,1) < Mg (2,8,2) < Al (2,8,3)
• m.pt: 98°C < 650°C < 659°C
• b.pt: 890°C < 1110°C < 2470°C
• Melting point increases across the period because the electron cloud density increases due to the greater number of electrons donated per atom. As a result, the ions are held more strongly.

Groups

• Li (2,1) > Na (2,8,1) > K (2,8,8,1)
• m.pt: 181°C > 98°C > 63°C
• b.pt: 1313°C > 890°C > 774°C
• Melting point decreases down a group.
• Ionic radius increases down the group. As the ions get bigger, the electron cloud becomes less effective holding them together, so they are easier to separate.

Properties of Metals

• High electrical conductivity: Valence electrons are highly mobile; voltage applied = flow of charge.
• High thermal conductivity: Delocalized electrons readily transfer energy; therefore, particles vibrate more rapidly.
• Malleable (bend) and ductile (stretched into wire).
• High MP and BP: Strong electrostatic bond increases melting and boiling point.
• High density.

Covalent Bonding

• Consists of a shared pair of electrons, with one electron being supplied by each atom either side of the bond.

Formation

• Between atoms of the same element: N2, O2, diamond, graphite
• Between atoms of different elements: CO2, SO2
• On the RHS of the table; when one of the elements is in theCCl4, SiCl_4
• Middle of the table; with head-of-the-group elements BeCl_2
• With high ionization energies

Hydrogen Molecule

• Each hydrogen atom needs one electron to complete its outer shell.
• Atoms share a pair of electrons to form a single covalent bond.

Hydrogen Chloride

• Hydrogen atom also needs one electron to complete its outer shell.
• Chlorine atom needs one electron to complete its outer shell.
• Atoms share a pair of electrons to form a single covalent bond.

Methane

• Each hydrogen atom needs 1 electron to complete its outer shell.
• A carbon atom needs 4 electrons to complete its outer shell.
• Carbon shares all 4 of its electrons to form 4 single covalent bonds.

Ammonia

• Each hydrogen atom needs one electron to complete its outer shell.
• Nitrogen atom needs 3 electrons to complete its outer shell.
• Nitrogen can only share 3 of its 5 electrons otherwise it will exceed the maximum of 8.
• A LONE PAIR REMAINS

Water

• Each hydrogen atom needs one electron to complete its outer shell.
• Oxygen atom needs 2 electrons to complete its outer shell.
• Oxygen can only share 2 of its 6 electrons otherwise it will exceed the maximum of 8.
• 2 LONE PAIRS REMAIN

Covalent Bonding in Methane

• Carbon (2.4) needs 4 electrons; hydrogen (1) needs 1 electron.
• Ratio of atoms:
  • element 1 4
  • CH_4

Covalent Bonding in Carbon Dioxide

• How do carbon and oxygen atoms form covalent bonds in a molecule of carbon dioxide?
  • Carbon (2.4) needs 4 electrons; oxygen (2.6) needs 2 electrons
  • Ratio of atoms: 1:2
  • CO_2 or O=C=O
    Double bonds: When two pairs of electrons are shared. In carbon dioxide there are two double bonds – one between each oxygen atom and the carbon atom. And show Lewis dot structure!

Dot and Cross Diagrams

• Hydrogen: Each atom needs one electron to complete its outer shell; atoms share a pair of electrons to form a single covalent bond
• Methane: Carbon needs four electrons to complete its outer shell; carbon shares all 4 of its electrons to form 4 single covalent bonds.
• Ammonia: Nitrogen needs three electrons to complete its outer shell; nitrogen can only share 3 of its 5 electrons otherwise it will exceed the maximum of 8; a lone pair remains.
• Water: Oxygen needs two electrons to complete its outer shell; oxygen can only share 2 of its 6 electrons otherwise it will exceed the maximum of 8; two lone pairs remain.
• Oxygen: Each atom needs two electrons to complete its outer shell; each oxygen shares 2 of its electrons to form a double covalent bond

Exceptions to the Octet Rule

• When there are an odd number of valence electrons
• When there are too few valence electrons
• When there are too many valence electrons

Simple Covalent Structures

• Covalent molecules that contain only a few atoms are called simple covalent structures.
• Most substances that contain simple covalent molecules have low melting and boiling points and are therefore liquids or gases at room temperature, e.g., water, oxygen, carbon dioxide, chlorine, and hydrogen.
• The covalent bonds within these molecules are strong, but the bonds between molecules are weak and easy to break.

Simple Covalent Molecules

• Bonding: Atoms are joined together within the molecule by covalent bonds.
• Electrical: Do not conduct electricity as they have no mobile ions or electrons.
• Solubility: Tend to be more soluble in organic solvents than in water; some are hydrolysed.
• Boiling point: Low - intermolecular forces are weak; they increase as molecules get a larger surface area (e.g., CH4 -161°C < C2H6 - 88°C < C3H8 -42°C).
• Because the intermolecular forces are weak, little energy is required to separate molecules from each other, so boiling points are low. Some boiling points are higher than expected for a given mass because you can get additional forces of attraction.

Structure of Molecular Solids

• A few substances that contain simple covalent molecules are solid at room temperature. These are molecular solids.
• The solid is formed because millions of iodine molecules are held together by weak forces of attraction to create a 3D molecular lattice.
• Two iodine atoms form a single covalent bond to become an iodine molecule.

Properties of Molecular Solids

• Low melting and boiling points
• Usually soft and brittle – they shatter when hit
• Cannot conduct electricity
• The weak forces of attraction between the molecules can be broken by a small amount of energy. This means that the molecular solids are soft and brittle and melt and boil at low temperatures.
• Molecular solids are also unable to conduct electricity because there are no free electrons or ions to carry a charge.

Molecular Solids: Iodine

• At room temperature and pressure, iodine is a greyish solid. However, it doesn’t need to be warmed much in order to produce a purple vapor.
• This is because iodine is composed of diatomic molecules (I_2) which exist in an ordered molecular crystal in the solid state.
• Each molecule is independent of the others, only being attracted by weak induced dipole-dipole interactions. Therefore, little energy is required to separate the iodine molecules.

Giant Covalent Structures

• In some substances, such as sand, diamond, and graphite, millions of atoms are joined together by covalent bonds.

• All the bonds are covalent, so giant covalent structures have very high melting and boiling points and are usually hard.

• The covalent bonds in these substances do not form molecules but vast networks of atoms called giant covalent structures.

• Many atoms joined together in a regular array by a large number of covalent bonds

Diamond and Silica

• Melting point: Very high because structure is made up of a large number of covalent bonds, all of which need to be broken if atoms are to be separated
• Electrical Conductivity: Don’t conduct electricity because they have no mobile ions or electrons, but graphite conducts electricity
• Strength: Hard because exists in a rigid tetrahedral structure. But graphite is soft

Diamond

• Melting Point: Very high because many covalent bonds must be broken to separate atoms.
• Strength: Strong because each carbon is joined to four others in a rigid structure.
• Electrical Conductivity: Non-conductor because no free electrons - all 4 carbon electrons used for bonding.

Graphite

• Melting Point: Very high because many covalent bonds must be broken to separate atoms
• Strength: Soft because each carbon is joined to three others in a layered structure and layers are held by weak induced dipole-dipole interactions that can slide over each other and used as a lubricant and in pencils
• Electrical Conductivity: Conductor because only three carbon electrons are used for bonding which leaves the fourth to move freely along layers layers can slide over each other used as a lubricant and in pencils

Structure of Sand

• Sand is mostly made of the mineral quartz, which is silicon dioxide.
• It has a giant covalent structure made up of silicon and oxygen atoms.
• Each silicon atom (2.8.4) is bonded to four oxygen atoms, and each oxygen atom (2.6) is bonded to two silicon atoms.

Silica

• Melting Point: Very high because many covalent bonds must be broken to separate atoms.
• Strength: Strong because each silicon atom is joined to four oxygens - each oxygen atom is joined to two silicons.
• Electrical Conductivity: Non-conductor - no mobile electrons.

Allotropes of Carbon

• Diamond and graphite appear to be very different substances, but what do they have in common?
• Both diamond and graphite are made up of carbon atoms.
• These allotropes of carbon have different properties because the atoms are bonded in different arrangements which create different giant structures.
• Different forms of the same element are called allotropes.

Carbon Compounds: Fullerenes

• A third class of carbon compounds have been discovered in recent years called fullerenes.
• Buckminsterfullerene is one type of fullerene. It contains 60 carbon atoms, each of which is bonded to three others by two single bonds and one double bond.
• The atoms in this allotrope of carbon form a sphere, like the shape of a football.
• The molecules can be called ‘bucky balls’. They are large but are not classified as giant structures.

Carbon Allotropes

• Nanotubes: Long hollow structures with walls
• Graphene:
  • Single layers of graphite
  • Strong and tough
  • Highly reactive carbon

Covalent Molecular Compound vs Network

Covalent Molecular Compound Properties

• Molecules are a group of atoms held together by covalent bonds.
• Between molecules, there are WEAK Intermolecular forces.
• Gives it a LOW melting point.
• Poor conductors of electricity.

Covalent Network Compound Properties

• Made of Covalent bonds.
• Lots of energy to break the bonds.
• Very high melting/boiling point.

Examples of Covalent Molecular Compound

• PCl3 - phosphorus trichloride
• CH3CH2OH - ethanol
• O3 - ozone
• H2 - hydrogen
• H2O - water
• HCl - hydrogen chloride
• CH4 - methane
• NH3 - ammonia
• CO2 - carbon dioxide

Examples of Covalent Network Compound

• Diamond
• graphite
• silicon dioxide
• silicon carbide

Bonding

Graphit

• Conducts

Covalent Molecules

• Insulators

Ionic Network

• Electrolytes

Metallic Network

• Conductors

Graphite and Diamond

• Q. Compare AND contrast between these two different structures

Bonding Affects Properties

• Q Does the type of bonding in a substance affect its properties?

Types of Bonding and Structures

• Giant ionic lattice: Millions of metal and non-metal ions, ionic solid .
• Simple molecular: Few non-metal atoms, usually liquid or solid, covalent.
• Giant covalent lattice: Millions of non-metal atoms, solid gigantic covalent
• Giant metallic lattice: Millions of metal ions, solid (except mercury – liquid), .metallic

Melting and Boiling Point

• Substances with giant structures generally have high melting and boiling points because all the atoms are strongly bonded together to form a continuous 3D lattice.
• A large amount of energy is needed to break these bonds.
  • Strong ionic bonds hold ions together
  • Strong covalent bonds hold atoms together
  • Strong metallic bonds hold ions together

Bonding Knowledge Check

• Recall the different types of physical and chemical bonding
• Understand how ionic, covalent, covalent, and metallic bonding arise
• Recall the different forms of covalent structures
• Understand how the physical properties depend on structure and bonding
• Understand how different types of physical bond have different strengths
• Construct diagrams to represent covalent bonding

Exam Question

• Iron is an extremely versatile material, used in many structures such as bridges, cars, ships, railway tracks and reinforced concrete. Iron(III) oxide is also known as the mineral hematite and is extracted from ores containing this mineral.
  • (a) Explain, in terms of structure and bonding, why iron is malleable yet iron(III) oxide is brittle.(6 marks)

Exam Marking Key

• Iron exhibits metallic bonding. or Iron exists as a sea of delocalized electrons surrounding positive metal cations. (1 mark)
• This bonding is non-directional in nature. (1 mark)
• Therefore when a force is applied the iron can change shape without disrupting the bonding. (1 mark)
• Iron(III) oxide exhibits ionic bonding. or Iron(III) oxide is composed of cations and anions. (1 mark)
• The ions are arranged in a rigid 3D lattice. (1 mark)
• Therefore when a force is applied the like charges align and repel, causing the substance to shatter. (1 mark)

Physical Properties of Diamond and Graphite

• Covalent Networks:
  • Diamond: Hard and strong = high melting point
  • Graphite: Slippery ( layers move relative to each other) due to freely-moving electrons between layers
  • Strong IMF forces in each later = high melting point

Allotropes

• Different forms of the same element are called allotropes.

Why Carbon Dioxide and Silicon Dioxide are Different

• Separation of mixtures:
  • Set 4 of STAWA : Mixtures
  • Set 13-17 + set19-20+ 24-25

Allotropes Exam Question

• The diamond company De Beers have used the phrase ‘A Diamond is Forever’ in their advertising campaign since 1948. However, interestingly, the conversion of diamond into graphite is known to occur via a spontaneous process, which can be represented as follows; C(s, diamond) \longrightarrow C(s, graphite). As can be seen in the chemical equation above, both forms of carbon have the same chemical formula.
  • (a) State the name given to structurally different forms of the same element. (1 mark)
  • Allotropes

High Melting Point Exam Question

• Both forms of carbon also have very high melting points.
  • (a) Explain, in terms of structure and bonding, why this is so. (3 marks)
• Answer: Both exhibit covalent network bonding. (1 mark) They consist of strong covalent bonds which need to be broken for melting to occur. (1 mark) Therefore, a high amount of heat / energy is required (resulting in a high melting point). (1 mark)

Exam Question

• There are, however, some properties which diamond and graphite do not share. For example, diamond is unable to conduct electricity whereas graphite can.
  • (c) Explain, in terms of structure and bonding, the difference in conductivity.(4 marks)
• Marking Key:
  • In diamond, each carbon atom is bonded to 4 other carbon atoms. (1 mark)
  • All electrons in diamond are therefore localized, leaving no mobile charge to conduct electricity. (1 mark)
  • In graphite, each carbon atom is bonded to 3 other carbon atoms. (1 mark)
  • The fourth valence electron of each carbon atom in graphite is delocalized, allowing it to conduct electricity. (1 mark)

Glossary

• allotrope – A structurally different form of an element with different physical properties.
• covalent bond – A strong bond between two atoms in which each atom shares one or more electrons with the other.
• covalent compound – A compound containing atoms joined by covalent bonds.
• double bond – A covalent bond in which each atom shares two of its electrons.
• giant structure – A structure containing millions of atoms or ions bonded together. The structure extends in three dimensions until all available atoms are used up.
• molecule – A small group of atoms which are held together by covalent bonds.
• molecular solid – A solid substance made up of molecules held together by weak forces of attraction, forming a lattice.
• single bond – A covalent bond in which each atom shares one of its electrons.
  • triple bond – A covalent bond in which each atom shares three of its electrons.