ATOMS, ELEMENTS AND COMPOUNDS

Some atoms are very reluctant to combine with other atoms and exist in the air around us as single atoms. These are the Noble Gases and have very stable electron arrangements e.g. 2, 2,8 and 2,8,8 because their outer shells are full. The first three are shown in the diagrams below and explains why Noble Gases are so reluctant to form compounds with other elements.

     (atomic number) electron arrangement

All other atoms therefore, bond together to become electronically more stable, that is to become like Noble Gases in electron arrangement. Bonding produces new substances and usually involves only the 'outer shell' or 'valency' electrons and atoms can bond in two ways.

The phrase CHEMICAL BOND refers to the strong electrical force of attraction between the atoms or ions in the structure. The combining power of an atom is sometimes referred to as its valency and its value is linked to the number of outer electrons of the original uncombined atom (see examples later).

Each type of chemical bonding is VERY briefly described below, with links to more detailed notes.

IONIC BONDING

Ionic bonds are formed by one atom transferring electrons to another atom to form ions.

Elements consist of neutral atoms or molecules, the electrical neutrality is because the number of positive protons equals the number of surrounding negative electrons in their respective energy levels.

Charged particles called IONS are atoms, or groups of atoms, which have lost or gained one or more electrons to have a overall net electrical positive charge or negative charge.

In losing or gaining electrons, the atoms try to attain a stable electron arrangement of a noble gas e.g. a full outer shell of electrons.

For a given atom, a nearly full shell will try to gain electrons and a nearly empty shell will tend to lose electrons

The atom losing electrons forms a positive ion (a cation) and is usually a metal.

The overall charge on the ion is positive due to excess positive nuclear charge (protons do NOT change in chemical reactions) e.g.

Group 1 alkali metals lose their single outer electron to form single positive ions e.g. Na ==> Na+ + 2e–

Group 2 metals lose their two outer electrons to form doubly charged positive ions e.g. Mg ==> Mg2+ + 2e–

The atom gaining electrons forms a negative ion (an anion) and is usually a non–metallic element.

The overall charge on the ion is negative because of the gain, and therefore excess, of negative electrons e.g.

Group 7 halogen atoms gain one electron to form a singly charged negative ion e.g. Cl + e– ==> Cl–

Group 6 non–metals gain two electrons to form a doubly charged negative ion e.g. O + 2e– ==> O2–

Therefore an IONIC BOND IS THE FORCE OF ATTRACTION BETWEEN ADJACENT IONS OF OPPOSITE CHARGE.

Which electronic structures are the most stable? because this what atoms will try to get to electronically!

     symbol (atomic number) electron arrangement

When atoms LOSE OR GAIN ELECTRONS, they try to attain the electron structure (electron configuration) of the electronically very stable atoms of the Group 0 Noble Gases eg helium (2), neon (2.8) or argon (2.8.8) etc. quite simply because these are the most stable electron arrangements with a full outer shell of electrons (full highest energy level).

In advanced level chemistry you will encounter examples of electronic structures of ions that are NOT those of a Noble Gas.

Its a good idea to have some idea of where the elements are in the periodic table, and their electronic structure, before looking at the theoretical electronic model for ionic bonding in

ionic compounds

The black zig–zag line 'roughly'

divides the metals on the left from the non–metals

on the right of the elements of the Periodic Table.

The electronic structures of the first 20 elements of the Periodic

Table

You need to know about these to understand the details of ionic

chemical bonding

mePart of the modern Period

metals => non–metals

ic Table

Pd = period, Gp = group

1H 

Note

that

H

does

not

read

ily fit

into

any

grou

p

at

o

mi

c

nu

m

be

r 

Chemical Symbol

eg 4

Transition MetalsGp 1 Alkal

i Metals  Gp 2 Alkaline Earth Metals  Gp 7

Halogens  Gp 0

Noble Gases

Chemica

l bo

nding comments about

the selected elements highlighted in white

e.g.

When

the metals

on the left combine with

the non–metals

on the right, an ionic bond is for

med

e.g.

the formation of an ionic compound like

sodium chloride NaCl

All the atoms of the metallic elements on the left–hand side of the Periodic Table eg (Groups 1/2) have 1/2 electrons in their outer shell (highest energy level) which are readily lost to form a positive ion of charge +1/+2 (cations) eg sodium, potassium, magnesium and calcium etc. The electronic structure of these stable positive ions are those of a Noble Gas with a full outer shell.

eg Na [2.8.1] ==> Na+ [2.8] like neon + e–, or Ca [2.8.8.2] ==> Ca2+ [2.8.8] like argon + 2e–

The atoms of the non–metallic elements on the right–hand side of the Periodic Table eg (Groups 6/7) have 6/7 electrons in their outer shell and try to gain 2/1 electrons to become electronically stable like a Noble Gas with a full outer shell of electrons eg oxygen and sulfur in Group 6 and the Group 7 Halogens – fluorine, chlorine, bromine and iodine. The electronic structure of these stable negative ions are those of a Noble Gas with a full outer shell.

eg O [2.6] + 2e– ==> O2– [2.8] like neon, or Cl [2.8.7] + e– ==> Cl– [2.8.8] like argon

Example 1: A Group 1 Alkali Metal + a Group 7 Halogen non–metal

e.g. sodium + chlorine ==> sodium chloride NaCl or ionic formula Na+Cl– 

In terms of electron arrangement in the formation of the ionic compound sodium chloride, the sodium donates its outer electron to a chlorine atom forming a single positive sodium ion and a single negative chloride ion.

The atoms have become stable ions, because electronically via electron transfer ...

... sodium becomes like neon (sodium ion, Na+) and chlorine like argon (chloride ion, Cl–).

Na (2.8.1) + Cl (2.8.7) ==> Na+ (2.8) Cl– (2.8.8)

can be summarised electronically to give the stable 'noble gas' structures as [2,8,1] + [2,8,7] ==> [2,8]+ [2,8,8]–

so both the sodium and chloride ions have a full outer shell like a noble gas

ONE  atom combines with ONE  atom to form 

Note in this electron diagram, only the original outer electrons are shown above.

The outer electron of the sodium atom (2.8.1) is transferred to the outer shell of the chlorine atom (2.8.7) giving it a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the same

time, the sodium ion also attains a stable noble gas electron structure (2.8).

The valencies of Na and Cl are both 1, that is, the numerical charge on the ions. sodium fluoride NaF, potassium bromide KBr and lithium iodide LiI etc. will all be electronically similar.

Only the outer valency electrons of the chloride ion are shown, the 'blob' electron represents the electron from the sodium atom which is accepted by the chlorine atom to form the chloride ion.

The charge on the sodium ion Na+ is +1 units (by convention shown as just +) because there is one more positive proton than there are negative electrons in the sodium ion (11p, 10e).

The charge on the chloride ion Cl– is –1 units (by convention shown as just –) because there is one more negative electron than there are positive protons in the chloride ion (17p, 18e).

Note:

would represent the full electronic structure diagram of the sodium ion [2.8] and the chloride ion [2.8.8], hence the full electronic structure of sodium chloride. Note that the 'blob' and 'x' electrons are identical, but their use is just a useful visual device to show how the ions are formed. The blue circle represents the nucleus.

The electronic dot & cross Lewis diagram for the ionic bonding in sodium chloride

Li is 2.1, K is 2.8.8.1, F is 2.7, rest of dot and cross diagram is up to you.

Gp1\7

F Cl Br I

Li LiF LiCl LiBr LiI

NaNaF

NaCl

NaBr

NaI

K KF KCl KBr KI

Rb RbRbC

lRbB

rRbI

CsCsF

CsCl

CsBr

CsI

All the formula highlighted in yellow can be described in the same way as sodium chloride

The Group 1 Alkali Metal atom loses one electron to form a singly charged positive ion

The Group 7 Halogen atom gains one electron to form a singly charged negative ion

Example 2: A Group 2 Alkaline Earth Metal + a Group 7 Halogen non–metal

e.g. magnesium + chlorine ==> magnesium chloride MgCl2 or ionic formula Mg2+(Cl–)2 

In terms of electron arrangement in the formation of the ionic compound magnesium chloride, the magnesium donates its two outer electrons to two chlorine atoms forming a double positive magnesium ion and two single negative chloride ions via electron transfer.

The atoms have become stable ions, because electronically, magnesium becomes like neon and chlorine like argon.

Mg (2.8.2) + 2Cl (2.8.7) ==> Mg2+ (2.8) 2Cl– (2.8.8)

can be summarised electronically as [2,8,2] + 2[2,8,7] ==> [2,8]2+ [2,8,8]–2  via electron transfer

so both the magnesium and chloride ions have a full outer shell of electrons like a noble gas

ONE  atom combines with TWO   atoms to

form 

Note in this electron diagram, only the original outer electrons are shown above.

The outer electrons of the magnesium atoms (2.8.2) are transferred to the outer shell of the chlorine atom (2.8.7) giving it a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the

same time, the magnesium ion also attains a stable noble gas electron structure (2.8).

NOTE

You can draw two separate chloride ions, but in these examples square brackets and a number subscript have been used, as in ordinary chemical formula.

The valency of Mg is 2 and chlorine 1, i.e. the numerical charges of the ions.

Beryllium fluoride BeF2, magnesium bromide MgBr2, calcium chloride CaCl2 or barium iodide BaI2 etc. will all be electronically similar.

represents the full electronic structure diagram of the magnesium ion [2.8] and the chloride ion [2.8.8], hence the full electronic structure of magnesium chloride.

Note that the 'blob' and 'x' electrons are identical, but their use is just a useful visual device to show how the ion is formed. The blue circle represents the nucleus.

The electronic dot & cross Lewis diagram for the ionic bonding in magnesium chloride

 

Ca is 2.8.8.2, Cl is 2.8.7, F is 2.7 rest of dot and cross diagrams are up to you, but calcium chloride is shown below.

The calcium atoms transfer their two outer electrons to the outer shell of two chlorine atoms

calcium chloride

The electronic dot & cross Lewis diagram for the ionic bonding in calcium chloride

The two outer electrons of the calcium atoms (2.8.8.2) are transferred to the outer shell of two chlorine atoms (2.8.7) until it has a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the same time, the calcium ion also attains a stable noble gas electron structure (2.8.8). The blue circle

represents the nucleus.

Gp2\7

F Cl Br I

Mg

MgF2

MgCl

2

MgBr2

MgI2

CaCaF2

CaCl

2

CaBr

2

CaI2

SrSrF2

SrCl

2

SrBr

2

SrI2

BaBaF2

BaCl

2

BaBr

2

BaI2

All the formula highlighted in yellow can be described in the same way as magnesium chloride or calcium chloride

The Group 2 Alkaline Earth Metal atom loses two electrons to form a doubly charged positive ion

The Group 7 Halogen atom gains one electron to form a singly charged negative ion

Example 3: A Group 3 metal + a Group 7 non–metal

e.g. aluminium + fluorine ==> aluminium fluoride AlF3 or ionic formula Al3+(F–)3 

In terms of electron arrangement in the formation of the ionic compound aluminium fluoride, the aluminium donates its three outer electrons to three fluorine atoms forming a triple positive aluminium ion and three single negative fluoride ions.

The atoms have become stable ions, because aluminium and fluorine becomes electronically like neon via electron transfer.

Valency of Al is 3 and F is 1, i.e. equal to the charges on the ions.

Al (2.8.3) + 3F (2.7) ==> Al3+ (2.8) 3F– (2.8)

can be summarised electronically as [2,8,3] + 3[2,7] ==> [2,8]3+ [2,8]–3

so both the aluminium and fluoride ions have a full outer shell like a noble gas

ONE  atom combines with THREE  atoms to

form 

Note in this electron diagram, only the original outer electrons are shown above.

The outer electrons of the aluminium atom (2.8.3) is transferred to the outer shell of the fluorine atoms (2.7) giving them a complete octet shell of outer electrons, just like a noble gas (2.8). At the same

time, the aluminium ion also attains a stable noble gas electron structure (2.8).

full electronic structure diagram of aluminium fluoride, the blue circle represents the nucleus.

The electronic dot & cross Lewis diagram for the ionic bonding in aluminium fluoride

Solid aluminium chloride/bromide/iodide have similar formula but are covalent when vapourised into Al2X6 dimer molecules, but AlCl3 has an ionic lattice in the solid, not sure on solid AlBr3 and AlI3, but these points are best left for an advanced AS–A2 chemistry discussion, not for GCSE students!

Example 4: A Group 1 Alkali Metal + a Group 6 non–metallic element

e.g. sodium/potassium + oxygen ==> sodium/potassium oxide Na2O/K2O or ionic formula (Na+)2O2–/(K+)2O2–

In terms of electron arrangement in the formation of the ionic compound sodium oxide, the two sodium/potassium atoms donate their outer electron to one oxygen atom.

This results in two single positive potassium ions to one double negative oxide ion via electron transfer.

All the ions have the stable electronic structures 2.8.8 (argon like) or 2.8 (neon like).

Valencies, sodium/potassium 1, oxygen/sulfur 2. giving the following formulae:

Lithium oxide, Li2O, sodium oxide Na2O, sodium sulphide Na2S and potassium K2S etc.

sodium oxide

2Na (2.8.1) + O (2.6) ==> 2Na+ (2.8.8) O2– (2.8)

can be summarised electronically as 2[2,8,1] + [2,6] ==> [2,8]+2 [2,8]2–

so both the sodium and oxide ions have a full outer shell like a noble gas

TWO  atoms combine with ONE  atom to form 

or

 +   +   ==> 

Note in this electron diagram, only the original outer electrons are shown above.

The outer electrons of the sodium atoms (2.8.1) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8). At the same time,

the sodium ion also attains a stable noble gas electron structure (2.8).

full electronic structure diagram of sodium oxide, the blue circle represents the nucleus.

The electronic dot & cross Lewis diagram for the ionic bonding in sodium oxide

 

potassium oxide

2K (2.8.8.1) + O (2.6) ==> 2K+ (2.8.8) O2– (2.8)

can be summarised electronically as 2[2,8,8,1] + [2,6] ==> [2,8,8]+2 [2,8]2–

so both the potassium and oxide ions have a full outer shell like a noble gas

TWO  atoms combine with ONE  atom to form 

Note in this electron diagram, only the original outer electrons are shown above.

The outer electrons of the potassium atoms (2.8.8.1) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8). At the same

time, the potassium ion also attains a stable noble gas electron structure (2.8.8).

full electronic structure diagram of potassium oxide,  the blue circle represents the nucleus.

The electronic dot & cross Lewis diagram for the ionic bonding in potassium oxide

 

The electronic similarities between the two examples are very obvious.

Li is 2.1, Na is 2.8.1, S is 2.8.6 (for group 1 sulphide compound), rest of dot and cross diagrams are up to you.

e.g. electronic structure diagrams for sodium sulfide Na2S and potassium sulfide K2S

sodium sulfide

The electronic dot & cross Lewis diagram for the ionic bonding in sodium sulphide

The outer electrons of the sodium atoms (2.8.1) are transferred to the outer shell of the sulfur atom (2.8.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the same

time, the sodium ion also attains a stable noble gas electron structure (2.8).

 

potassium sulfide

The outer electrons of the potassium atoms (2.8.8.1) are transferred to the outer shell of the sulfur atom (2.8.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the

same time, the potassium ion also attains a stable noble gas electron structure (2.8.8).

The electronic dot & cross Lewis diagram for the ionic bonding in potassium sulphide

Gp1\6

O S

Li Li2O Li2S

NaN

a2ONa2S

K K2O K2S

RbR

b2OR

b2S

CsC

s2OCs2S

All the formula highlighted in yellow can be described in the same way as sodium oxide, potassium oxide, sodium sulfide or calcium sulphide

The Group 1 Alkali Metal atom loses one electron to form a singly charged positive ion

The Group 6 non–metal atom gains two electrons to form a doubly charged negative ion

Example 5: A Group 2 Alkaline Earth Metal + a Group 6 non–metallic element

e.g. magnesium/calcium + oxygen ==> magnesium/calcium oxide MgO/CaO or ionic formula Mg2+O2–/Ca2+O2–

In terms of electron arrangement in the formation of the ionic compound magnesium oxide, one magnesium/calcium atom donates its two outer electrons to one oxygen atom.

This results in a double positive calcium ion to one double negative oxide ion via electron transfer.

All the ions have the stable electronic structures 2.8.8 (argon like) or 2.8 (neon like). the valency of both calcium and oxygen is 2.

magnesium oxide

ONE  atom combines with ONE  atom to

form 

Note in this electron diagram, only the original outer electrons are shown above.

The two outer electrons of the magnesium atoms (2.8.2) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8). At

the same time, the magnesium ion also attains a stable noble gas electron structure (2.8).

full electronic structure of magnesium oxide

For magnesium oxide: Mg (2.8.2) + O (2.6) ==> Mg2+ (2.8) O2– (2.8)

the stable 'noble gas' structures can be summarised electronically as [2,8] + [2,6] ==> [2,8]2+ [2,8]2–

so both the magnesium and oxide ions have a full outer shell like a noble gas

The electronic dot & cross Lewis diagram for the ionic bonding in magnesium oxide

 

calcium oxide

Ca (2.8.8.2) + O (2.6) ==> Ca2+ (2.8.8) O2– (2.8)

can be summarised electronically as [2,8,8,2] + [2,6] ==> [2,8,8]2+ [2,8]2–

ONE  atom combines with ONE  atom to

form 

Note in this electron diagram, only the original outer electrons are shown above.

The two outer electrons of the calcium atoms (2.8.8.2) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8). At the same

time, the calcium ion also attains a stable noble gas electron structure (2.8.8).

full electronic structure of calcium oxide

The electronic dot & cross Lewis diagram for the ionic bonding in calcium oxide

 

Magnesium oxide MgO, magnesium sulphide MgS and calcium sulphide CaS will be similar electronically and give identical giant ionic lattice structures.

Group 2 metals lose the two outer electrons to give the stable 2+ positive ion (cation) and S and O, both non–metals in Group 6, have 6 outer electrons and gain 2 electrons to form 2– negative ion (anion).

For magnesium sulphide: Mg (2.8.2) + S (2.8.6) ==> Mg2+ (2.8) S2– (2.8.8)

For calcium sulphide: Ca (2.8.8.2) + S (2.8.6) ==> Ca2+ (2.8.8) S2– (2.8.8)

so both the magnesium/calcium and sulfide ions have a full outer shell like a noble gas

The dot and cross (ox) diagrams will be identical to that for calcium oxide above, except Mg instead of Ca (same group) and S instead of O (same group of Periodic Table). eg

electronic structure of magnesium sulfide MgS

electronic structure of calcium sulfide CaS

The electronic dot & cross Lewis diagrams for the ionic bonding in magnesium sulphide and calcium sulphide

Gp2\6 O S

Mg MgO MgS

Ca CaO CaS

Sr SrO SrS

Ba BaO BaS

All the formula highlighted in yellow can be described in the same way as magnesium oxide, magnesium sulphide, calcium oxide or calcium sulphide

The Group 2 Alkaline Earth Metal atom loses two electrons to form a doubly charged positive ion

The Group 6 non–metal atom gains two electrons to form a doubly charged negative ion

Example 6: A Group 3 metal + a Group 6 non–metal

e.g. aluminium + oxygen ==> aluminium oxide Al2O3 or ionic formula (Al3+)2(O2–)3

In terms of electron arrangement in the formation of the ionic compound aluminium oxide, two aluminium atoms donate their three outer electrons to three oxygen atoms.

This results in two triple positive aluminium ions to three double negative oxide ions via electron transfer.

All the ions have the stable electronic structure of neon 2.8. Valencies, Al = 3 and O = 2

2Al (2.8.3) + 3O (2.6) ==> 2Al3+ (2.8) 3O2– (2.8)

can be summarised electronically as 2[2,8,3] + 3[2,6] ==> [2,8]3+2 [2,8]2–

3

so both the aluminium and oxide ions have a full outer shell like a noble gas

TWO   atoms combine with THREE   atoms to

form 

Note in this electron diagram, only the original outer electrons are shown above.

The three outer electrons of the aluminium atoms (2.8.3) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas 2.8). At

the same time, the aluminium ion also attains a stable noble gas electron structure (2.8).

Note:

The charge on the aluminium ion Al3+ is +3 units (shown as 3+) because there are three more positive protons than there are negative electrons in the aluminium ion.

The charge on the oxide ion O2– is –2 units (shown as 2–) because there are two more negative electrons than there are positive protons in the oxide ion.

full electronic structure of aluminium oxide

A GIANT IONIC LATTICE – explaining its properties

The diagram on the right is typical of the giant ionic crystal structure of ionic compounds like sodium chloride and magnesium oxide.

Solid ionic compounds consist of a giant lattice of closely packed ions which are all combine together to form a crystal.

The alternate positive and negative ions in an ionic solid are arranged in an orderly/regular way in a giant ionic lattice structure shown on the right.

The ionic bond is the strong electrical attraction between the oppositely charged positive and negative ions next to each other in the lattice.

The bonding extends throughout the crystal in all directions. Salts and metal oxides are typical ionic compounds. This strong bonding force makes the structure hard (if brittle) and have high melting and

very high boiling points, so they are not very volatile! A relatively large amount of energy is needed to melt or boil ionic compounds to

reduce/overcome the strong bonding forces.o Energy changes for the physical changes of state of melting and boiling for a range

of differently bonded substances are compared in a section of the Energetics Notes.

The bigger the charges on the ions the stronger the bonding attraction e.g. magnesium oxideMg2+O2– has a much higher melting point than sodium chloride Na+Cl–.

o The ions of magnesium oxide are both doubly charged so the electrostatic attraction is much greater (its actually about 4x as strong attractive force).

As it happens in this case, the ions in magnesium oxide are smaller than the ions in sodium chloride, so the ions in magnesium oxide can pack closer together and this also increase the attractive bonding force.

o This double effect results in a much stronger ionic bond in magnesium oxide, so a much greater thermal kinetic energy i.e. a much greater temperature, is required to weaken the giant ionic lattice and melt the crystals of magnesium oxide compared to sodium chloride.

o Simple experimental evidence – sodium chloride melts at 801oC, whereas magnesium oxide melts much higher at 2852oC.

Unlike covalent molecules, ALL ionic compounds are crystalline solids at room temperature.

They are hard but brittle, when stressed the bonds are broken along planes of ions which shear away.

o They are NOT malleable like metals. Many ionic compounds are soluble in water, but not all, so don't make this assumption.

o Salts can dissolve in water because the ions can separate and become surrounded by water molecules which weakly bond to the ions (see diagrams below).

o This reduces the attractive forces between the ions, preventing the crystal structure to exist.

o Evaporating the water from a salt solution will eventually allow the ionic crystal lattice to reform.

The solid crystals DO NOT conduct electricity because the ions are not free to move to carry an electric current.

o However, if the ionic compound is melted or dissolved in water, the liquid will now conduct electricity, as the ion particles are now free to move and carry the electric current in the molten salt or the solution of the salt in aqueous solution (see diagrams below).

o This electrical conduction under these conditions is evidence for the existence of ions in this type of compound.

 

An 'advanced' molecular particle picture of sodium chloride dissolving in water

(the partial electrical charges δ+ and δ– are for advanced level students only)

 

 ==> 

solid sodium chloride ==> molten sodium chloride (fixed ions to free moving ions)

How to work out the formula for an ionic compound

Table 1a. selected ions and charges

Table 1b. The periodic table pattern of charges on ions

The charge is based on the number of electrons lost (giving positive ions) or gained (giving negative ions) forming a noble gas electron structure, i.e. to make a full outer shell of electrons

CATIONS from metals ANIONS from non–metals

Group 1 Group 2 Group 3 Group 6 Group 7

lithium ion Li+ beryllium ion Be2+   oxide ion O2– fluoride ion F–

sodium ion Na+ magnesium ion Mg2+ aluminium ion Al3+ sulfide ion S2– chloride Cl–

potassium ion K+ calcium ion Ca2+     bromide ion Br–

        iodide ion I–

COVALENT BOND – electron sharing in big or small molecule

Covalent bonds are formed by atoms sharing electrons to form bonds that hold the atoms together in a molecule.

This type of bond usually formed between two non–metallic elements. The molecules might be that of an element i.e. one type of atom only OR from different elements chemically combined to form a compound.

Note: The molecular formula is the summary of all the atoms in a molecule.

The covalent bonding is caused by the mutual electrical attraction between the two positive nuclei of the two atoms of the bond, and the SHARING the negative electrons between them.

A COVALENT BOND IS THE SHARING OF ELECTRONS BETWEEN TWO ATOMS

It only involves electrons in the outer shell i.e. the outermost energy level containing 1–7 electrons, which can be shared between atoms to form a covalent bond.

One single covalent bond is a sharing of 1 pair of electrons, two pairs of shared electrons between the same two atoms gives a double bond and it is possible for two atoms to share 3

pairs of electrons and give a triple bond.

Note: In the examples of covalent bonding it is assumed YOU can work out the electron configuration (arrangement in shells or energy levels) given the atomic number from the Periodic Table.

This kind of bond or electronic linkage does act in a particular direction i.e. along the 'line' between the two nuclei of the atoms bonded together, this is why covalent molecules have a particular shape.

In the case of ionic or metallic bonding, the electrical attractive forces act in all directions around the particles involved.

Which electronic structures are the most stable? because is this what atoms will try to get to electronically!

     symbol (atomic number) electron arrangement

When atoms SHARE ELECTRONS in a covalent bond, they try to attain the electron structure (electron configuration) of the electronically very stable atoms of the Group 0 Noble Gases eg helium (2), neon (2.8) or argon (2.8.8), that is, a full outer shell of electrons (full highest energy level).

Quite simply, this is because these are the most stable electron arrangements and have a full outer shell of electrons (full highest energy level).

The number of bonds formed depends on the number of electrons that needs to be shared so that any pair of atoms in a molecule forming a covalent bond attain the electron arrangement of a noble gas (i.e. 2, 2.8 or 2.8.8 etc.)

Note that hydrogen and helium only have one shell, so when referring to the full outer shell of hydrogen, it is the one and only shell, but the descriptive word 'outer' is much more crucial when describing the electronic structures of any element with at least two shells e.g. when describing covalent bonding in molecules containing carbon, oxygen, nitrogen and chlorine etc.

In advanced level chemistry you will encounter examples of electronic structures of atoms in covalent molecules that are NOT those of a Noble Gas.

Its a good idea to have some idea of where the elements are in the periodic table, and their electronic structure, before looking at the theoretical electronic model for covalent bonding

in molecules (elements or compounds).

The black zig–zag line 'roughly' divides the

metals on the left from the non–metals on the

right of the elements of the Periodic Table.

The electronic structures of the first 20 elements of the Periodic

Table

You need to know about these to understand the

details of covalent chemical

bonding

me

Part of the modern Periodic Table

Pd = period, Gp = group

metals => non–metals

1H 

Note

that

H

does

not

read

ily fit

into

any

grou

p

at

o

mi

c

n

u

m

be

r 

Chemical Symbol eg 4

Be

Transition MetalsGp 1 Alkal

i Metals  Gp 2 Alkal

ine Earth Metals  Gp 7

Halogens  Gp 0

Noble Gases

Chemica

l bonding comments about

the selected elements highlighted in white

When any

two of

the

highlighted non–metals

on the right of

the periodic table

(and hydrogen) combine with each other

OR with

themselves, covalent bonds

are formed, e.g

. the forma

tion of a

covalent compounds

like

hydrogen chloride HCl,

sulfur dioxide S

O2, and

element

molecules

like

hydrogen H2 or oxygen O2.

Example 1: Covalent bonding diagram for HYDROGEN covalent molecule, molecular formula H2

Two hydrogen atoms (1) form the molecule of the element hydrogen H2

Hydrogen H2 is one short of a full shell like helium, so two hydrogen atoms share each others electron to have a full outer shell.

 and   combine to form   where both atoms have a pseudo helium structure of 2 outer electrons around each atom's nucleus. Any covalent bond (like H–H) is formed from the mutual attraction of two positive nuclei and negative electrons between them (i.e. effectively 'electron sharing'). The nuclei would be a tiny dot in the middle of where the H symbols are drawn! H valency is 1.

 simplified 'dot and cross' electronic diagram for the covalently bonded hydrogen molecule

The hydrogen molecule is held together by the strong hydrogen–hydrogen single covalent bond H–H (displayed formula)..

Remember, electronically, hydrogen is simply 1 and becomes like helium 2, so the hydrogen atoms effectively have a full outer shell in forming the covalent bonds when the atoms share their outer electrons.

Example 2:  Covalent bonding diagram for CHLORINE covalent molecule, molecular formula Cl2

Two chlorine atoms (2.8.7) form the molecule of the element chlorine Cl2

Chlorine Cl2 is one electron short of a full outer shell of 8 like argon, so two chlorine atoms share an electron to have full outer shells.

 and   combine to form   where both atoms have a pseudo argon structure of 8 outer electrons around each atom. The electronic dot & cross Lewis diagrams for covalent bonding in chlorine

 simplified 'dot and cross' electronic diagram for the covalently bonded chlorine molecule

The chlorine molecule is held together by the strong chlorine–chlorine single covalent bond from sharing outer electrons,  Cl–Cl (displayed formula)..

Electronically, both chlorines (2.8.7) become like argon (2.8.8), so the chlorine atoms effectively have a full outer shell in forming the covalent bonds when the atoms share their outer electrons.

All the other halogens would be similar e.g. F2, Br2 and I2 etc.

Here the valency of halogens like chlorine is 1.

Note that the two inner shells of chlorine's electrons are not shown in the diagram above, as they are in the diagram on the left, and, remember, 'blobs' and 'crosses' are all the same electrons in their specific energy levels and only the outer shells of electrons are involved in the covalent bonding here.

is the full 'dot and cross' electronic diagram for the

covalent bonding in the chlorine molecule.

Example 3:  Covalent bonding diagram for HYDROGEN CHLORIDE covalent molecule, molecular formula HCl

One atom of hydrogen (1) combines with one atom of chlorine (2.8.7) to form the molecule of the compound hydrogen chloride HCl

Both hydrogen and chlorine have one electron short of a full outer shell (2 for H, 8 for Cl), so both atoms share an electron to have full outer shells.

 and   combine to form   where hydrogen is electronically like helium (2) and chlorine like argon (2.8.8).

The hydrogen chloride molecule is held together by the strong hydrogen–chlorine single covalent bond by sharing electrons, H–Cl (displayed formula).

Note that the two inner shells of chlorine's electrons (2.8.7) are NOT shown (see chlorine atom diagram in example 2.

Electronically, hydrogen (1) becomes like helium (2) and chlorine (2.8.7) becomes like argon (2.8.8), so the hydrogen and chlorine atoms effectively have full outer shells in forming the covalent bonds when the atoms share their outer electrons. The two inner shells of chlorine's electrons are not shown, only the outer shells of electrons are involved in the covalent bonding here.

 simplified 'dot and cross' electronic diagram for the covalently bonded hydrogen chloride molecule

is the full 'dot and cross' electronic diagram for the covalent bonding in the hydrogen chloride molecule.

The electronic dot & cross Lewis diagrams for covalent bonding in hydrogen chloride.

All the other hydrogen halides will be similar e.g. hydrogen fluoride HF, hydrogen bromide HBr and hydrogen iodide HI.

Note: Hydrogen chloride gas is a true covalent substance consisting of small HCl molecules. If the gas is dissolved in a hydrocarbon solvent like hexane or methylbenzene it remains as HCl molecules and because there are no ions present, the solution does not conduct electricity. However, if hydrogen chloride gas is dissolved in water, things are very different and the HCl molecules split into ions. Hydrochloric acid is formed which consists of a solution of hydrogen ions (H+) and chloride ions (Cl–). The solution then conducts electricity and passage of a d.c.

current causes electrolysis to take place forming hydrogen and chlorine.

Example 4:  Covalent bonding diagram for WATER covalent molecule, molecular formula H2O

Two atoms of hydrogen (1) combine with one atom of oxygen (2.6) to form the molecule of the compound water H2O

Hydrogen is one electron short of a full shell, oxygen is two electrons short of a full outer shell of 8, so two hydrogen atoms share their electrons with the six outer electrons of oxygen, so all three atoms now have a full outer shell.

 and  and   combine to form   so that the hydrogen atoms are electronically like helium and the oxygen atom becomes like neon (2.8, but only the outer shell of oxygen's electrons are shown).

Electronically, hydrogen (1) becomes like helium (2) and oxygen (2.6) becomes like neon (2.8), so the hydrogen and oxygen atoms effectively have full outer shells in forming the covalent bonds when the atoms share their outer electrons.

 simplified 'dot and cross' electronic diagram for the covalently bonded water molecule

The water molecule is held together by the strong H–O hydrogen–oxygen single covalent bonds by sharing electrons.

Note that the inner shell of oxygen's electrons are shown in this diagram, but not in the bonding diagram. Only the outer shell of oxygen's electrons are involved in the covalent bonding here.

The molecule can be shown as  (displayed formula) with two hydrogen – oxygen single covalent bonds (AS note: called a V or bent shape, the H–O–H bond angle is 105o). The two pairs of double dots represent pairs of electrons not involved in the covalent bonding in water. Hydrogen sulphide will be similar, since sulphur (2.8.6) is in the same Group 6 as oxygen. Valency of oxygen and sulphur is 2 here.

is the full 'dot and cross' electronic diagram for the covalent bonding in the water molecule.

The electronic dot & cross Lewis diagrams for covalent bonding in water

Example 5:  Covalent bonding diagram for AMMONIA covalent molecule, molecular formula NH3

Three atoms of hydrogen (1) combine with one atom of nitrogen (2.5) to form the molecule of the compound ammonia NH3

Each hydrogen atom is one electron short of a helium structure (full shell) and nitrogen is three electrons short of a full outer shell (of 8), so three hydrogen atoms share their electrons with the five outer electrons of nitrogen, so all four atoms effectively have full outer shells.

three of   and one   combine to form   so that the hydrogen atoms are electronically like helium and the nitrogen atom becomes like neon (only the outer shell of nitrogen's electrons are shown).

Electronically, hydrogen (1) becomes like helium (2) and nitrogen (2.5) becomes like neon (2.8), so the hydrogen and nitrogen atoms effectively have full outer shells in forming the covalent bonds when the atoms share their outer electrons.

simplified 'dot and cross' electronic diagram for the covalently bonded ammonia molecule

The ammonia molecule is held together by the strong N–H nitrogen–hydrogen single covalent bonds by sharing electrons.

Note that the inner shell of nitrogen's electrons are not shown (as in this diagram on the right), only the outer shell of nitrogen's electrons are involved in the covalent bonding here.

The molecule can be shown as   (displayed formula) with three nitrogen – hydrogen single covalent bonds (AS note: called a trigonal pyramid shape, the H–N–H bond angle is 107o). The double dots represent a pair of electrons not involved in the covalent bonding in ammonia. PH3 will be similar since phosphorus (2.8.5) is in the same Group 5 as nitrogen. Valency of nitrogen or phosphorus is 3 here.

is the full 'dot and cross' electronic diagram for the covalent bonding in the ammonia molecule.

The electronic dot & cross Lewis diagrams for covalent bonding in ammonia

Example 6: Covalent bonding diagram for METHANE covalent molecule, molecular formula CH4

Four atoms of hydrogen (1) combine with one atom of carbon (2.4) to form the molecule of the compound methane CH4

Each hydrogen atom is one electron short of a helium structure (full shell) and carbon is four electrons short of a full outer shell (of 8), so four hydrogen atoms share their electrons with the four outer electrons of carbon, so all five atoms effectively have full outer shells.

four of   and one of   combine to form   so that the hydrogen atoms are electronically like helium and the carbon atom becomes like neon (only the outer shell of

carbon's bonding electrons are shown).

Electronically, hydrogen (1) becomes like helium (2) and carbon (2.4) becomes like neon (2.8), so the hydrogen and carbon atoms effectively have full outer shells in forming the covalent bonds when the atoms share their outer electrons.

 simplified 'dot and cross' electronic diagram for the covalently bonded methane molecule

The methane molecule is held together by the four strong C–H carbon–hydrogen covalent bonds by sharing electrons.

Note that the inner shell of carbon's electrons are not shown above, only the outer shell of carbon's electrons are involved in the covalent bonding.

The molecule can be shown as   (displayed formula) with four carbon – hydrogen single covalent bonds (AS note: called a tetrahedral shape, the H–C–H bond angle is 109o). SiH4 will be similar because silicon (2.8.4) is in the same group as carbon.

All the bonds in the above examples are single covalent bonds. Below are three examples 7–9, where there is a double bond in the molecule, in order that the atoms have stable Noble Gas outer electron arrangements around each atom. Carbon and silicon have a valency of 4.

More complex examples can be worked out e.g. involving C, H and O. In each case link in the atoms so that there are 2 around a H (electronically like He), or 8 around the C or O (electronically like Ne).

is the full 'dot and cross' electronic diagram for the covalent bonding in the methane molecule.

The electronic dot & cross Lewis diagrams for covalent bonding in methane

Example 7:  Covalent bonding diagram for OXYGEN covalent molecule, molecular formula O2

 Two atoms of oxygen (2.6) combine to form the molecules of the element oxygen O2 (only the outer shell of oxygen's electrons are shown).

Each oxygen atom is two electrons short of a full outer shell, so each oxygen atom shares two of its electrons with the other atom, so both oxygen atoms have a full outer shell.

 simplified 'dot and cross' electronic diagram for the covalently bonded oxygen molecule

The molecule has one O=O double covalent bond  (displayed formula), oxygen valency 2.

The oxygen molecule is held together by the strong O=O oxygen–oxygen double covalent bond by sharing electrons.

Electronically, by sharing two electrons, both oxygen atoms attain a pseudo neon structure (2.8), so the oxygen atoms effectively have full outer shells in forming the covalent bonds when the atoms share their outer electrons.

is the full 'dot and cross' electronic diagram for the covalent bonding in the oxygen molecule.

The electronic dot & cross Lewis diagrams for covalent bonding in oxygen

Example 8: Covalent bonding diagram for CARBON DIOXIDE covalent molecule, molecular formula CO2

   One atom of carbon (2.4) combines with two atoms of oxygen (2.6) to form the compound carbon dioxide CO2 (only the outer shell of carbon's electrons are shown).

Carbon is four electrons short of a full outer shell (8 electrons) and oxygen is two electrons short of

a full outer shell (8 electrons), so one carbon atom shares its four outer electrons with two outer electrons from each of the oxygen atoms, so all three atoms now have a full outer shell of 8 electrons in the formation of two double bonds (O=C=O).

Electronically, carbon (2.4) becomes like neon (2.8) and oxygen (2.6), also becomes like neon (2.8), so the hydrogen and carbon atoms effectively have full outer shells in forming the covalent bonds when the atoms share their outer electrons.

 simplified 'dot and cross' electronic diagram for the covalently bonded carbon dioxide molecule

is the full 'dot and cross' electronic diagram for the covalent bonding in the carbon dioxide molecule.

The electronic dot & cross Lewis diagrams for covalent bonding in carbon dioxide

The molecule can be shown as   (displayed formula) with two carbon = oxygen double covalent bonds (AS note: called a linear shape, the O=C=O bond angle is 180o). Valencies of C and O are 4 and 2 respectively.

The carbon dioxide molecule is held together by the strong C=O carbon–oxygen double covalent bonds by sharing electrons.

Example 9: Covalent bonding diagram for ETHENE covalent molecule, molecular formula C2H4

 Two atoms of carbon (2.4) combine with four atoms of hydrogen (1) to form ethene C2H4 (only the outer shell of carbon's electrons are shown).

 simplified 'dot and cross' electronic diagram for the covalently bonded ethene molecule

Electronically, hydrogen (1) becomes like helium (2) and carbon (2.4) becomes like neon (2.8), so ALL the hydrogen and carbon atoms effectively have full outer shells in forming the covalent bonds

when the atoms share their outer electrons.

The molecule can be shown as   (displayed formula) with one carbon = carbon double bond and four carbon – hydrogen single covalent bonds (called aplanar shape, its completely flat!, the H–C=C and H–C–H bond angles are 120o). The valency of carbon is still 4.

The ethane molecule is held together by the four strong C–H carbon–hydrogen single covalent

bonds and one C=C carbon–carbon covalent double bond.

Typical properties of simple molecular substances

Composed of relatively small covalently bonded molecules

Why do simple covalent molecules typically have low melting and low boiling points?o Typical examples are water (ice & liquid), petrol, butter etc.o The particles in the above diagram represent whole molecules, but the general

picture of particles in the three states of matter help to understand the properties of simple molecular substances.

Why do simple molecules NOT usually conduct electricity even when liquid/molten/dissolved.

The first point to appreciate is that the chemical bonding forces between atoms in a molecule are strong BUT the bonding forces between small simple molecules are weak. These weak electrical attractive forces are known as 'intermolecular forces' or 'intermolecular bonding'.

o DO NOT CONFUSE THESE TWO FORCES or it makes the following discussion on the physical properties of simple molecules difficult to follow.

o The contrast between the strong bonds between atoms in a molecule and the weak bonds between individual molecules is really important to know understand the consequences.

o It will also help you to understand why covalent giant molecular structures have very different physical properties.

The electrical forces of attraction, that is the chemical bond, between atoms in a molecule are usually very strong, so,  most covalent molecules do not change chemically on moderate heating.

o e.g. although a covalent molecule like iodine, I2, is readily vapourised on heating, it does NOT break up into iodine atoms I. The purple vapour you see on heating iodine is entirely composed of the diatomic I2 molecules.

o The I–I covalent bond is strong enough to withstand the heating and so the purple vapour still consists of the same I2 molecules as the dark coloured solid is made up of.

o In other words, on heating a simple molecular material like iodine, heating weakens

the forces between the molecules BUT not the forces between the atoms in the molecule.

Chemical bonds between atoms are generally only broken if a substance is heated to a VERY high temperature like in the cracking break–down reactions of alkanes from crude oil.

So why the ease of vaporisation on heating?o Although the bonding between the atoms within a molecules is very strong the

electrical attractive force between individual molecules is very weak, so the bulk material is not very strong physically and this has consequences for the melting points and boiling points.

oo If you take the Group 7 Halogen molecules, the F–F, Cl–Cl, Br–Br, I–I covalent

bonds (–) are very strong, but the F–F....F–F, Cl–Cl....Cl–Cl, Br–Br....Br–Br and I–I....I–I intermolecular

bonds are weak (weak intermolecular bonding shown as ....), resulting in low melting/boiling points e.g. at room temperature fluorine and

chlorine are gases, bromine a low boiling liquid and iodine an easily vapourised solid on gentle heating.

Note: The bigger the molecule, the stronger the intermolecular forces, which is why the melting/boiling points increase down group 7.

Similarly for hydrocarbons like alkanes, the longer the molecule, the higher the boiling point.

These weak electrical attractions are known as intermolecular forces (or intermolecular bonding) and are readily weakened further on heating. In a solid, the effect of absorbing heat energy results in increased the thermal vibration of the molecules which weakens the intermolecular forces. In liquids the increase in the average particle kinetic energy makes it easier for molecules to overcome the intermolecular forces and change into a gas or vapour.  Consequently, small covalent molecules tend to be volatile liquids with low boiling points, so easily vapourised, or low melting point solids.

o So, on heating simple molecular substances (small molecules) the inter–molecular forces are easily overcome with the increased kinetic energy of the particles, giving the material a low melting or boiling point because a relatively low value of kinetic energy is needed to effect these state changes.

o Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances.

o This contrasts with the high melting points of giant covalent structures with their strong 3D network.

o Reminder: The weak electrical attractive forces between molecules, the so called intermolecular forces should be clearly distinguished between the strong covalent bonding between atoms in molecules (small or giant), and these are sometimes referred to as intramolecular forces (i.e. internal to the molecule).

Covalent structures are usually poor conductors of electricity because there are no free

electrons or ions in any state to carry electric charge. Most small molecules will dissolve in some solvent to form a solution.

o Hydrocarbon molecules like hexane or paraffin wax dissolve in organic solvents but not water, but sugars are also low melting small covalent molecules but do dissolve in water, insoluble in hydrocarbon solvents.

The properties of these simple small molecules should be compared and contrasted with those molecules of a giant covalent nature (next section).

o Apart from points on the strong bonds between the atoms in the molecule and the lack of electrical conduction, all the other properties are significantly different!

COVALENT BONDING – macromolecules & giant covalent structures

The structure of the three allotropes of carbon (diamond, graphite and fullerenes), silicon and silicon dioxide (silica)

DIAGRAMS

It is possible for many atoms to link up to form a giant covalent structure or lattice.

o The structures of giant covalent structure are usually based on non–metal atoms like carbon, silicon and boron.

o The atoms in a giant covalent lattice are held together by strong directional covalent bonds and every atoms is connected to at least 2, 3 or 4 atoms.

o What you might call 'atomic networking'! This very strong 3–dimensional covalent bond network or

lattice gives the structure great thermal stability e.g. very high melting point and often great physical strength.

o This is because it takes so much thermal kinetic energy to weaken the bonds sufficiently to allow melting.

This gives them significantly different properties from the small simple covalent molecules (see simple molecular substances).

This is illustrated by carbon in the form of diamond (an allotrope of carbon). Carbon has four outer electrons that form four single bonds, so each carbon bonds to four others by electron pairing/sharing.

o Pure silicon, another element in Group 4, has a similar structure.

o NOTE: Allotropes are different forms of the same element in the same physical state. They occur due to different bonding arrangements and so diamond, graphite (below) and fullerenes (below) are the three solid allotropes of the element carbon.

Oxygen (dioxygen), O2, and ozone (trioxygen), O3, are the two small gaseous allotrope molecules of the element oxygen.

Sulphur has three solid allotropes, two different crystalline forms based on small S8 molecules called rhombic and monoclinic sulphur and a 3rd form of long chain ( –S–S–S– etc.) molecules called plastic sulphur.

TYPICAL PROPERTIES of GIANT COVALENT STRUCTURES

This type of giant covalent structure is thermally very stable

DIAMOND

  

 

 

  SILICA

silicon dioxide

and has a very high melting and boiling points because of the strong covalent bond network (3D or 2D in the case of graphite below).

A relatively large amount of energy is needed to melt or boil giant covalent structures because strong chemical bonds must be broken (and not just weakening intermolecular forces as in the case of small covalent molecules like water).

o Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances are compared in a section of the Energetics Notes.

They are usually poor conductors of electricity because the electrons are not usually free to move as they are in metallic structures.

o All the valency bonding electrons are tightly held and shared by the two atoms of any bond, so in giant covalent structures they are rarely free to move through the lattice and not even when molten either, since these giant molecular covalent structures do NOT contain ions.

Also, because of the strength of the bonding in all directions in the structure, they are often very hard, strong and will not dissolvein solvents like water. The bonding network is too strong to allow the atoms to become surrounded by solvent molecules

Silicon dioxide (silica, SiO2) has a similar 3D structure and properties to carbon (diamond) shown below and also pure silicon itself.

DIAMOND Cn where n is an extremely large number of carbon atoms!

o In diamond every carbon atoms is strongly linked to four other carbon atoms by strong directional covalent bonds giving a very strong lattice.

Theoretically in a diamond crystal all the carbon atoms are linked together.

The result is a very pure crystal structure with a high refractive index that gives diamonds quite a sparkle as light passes through it.

o The hardness of carbon in the form of diamond enables it to be used as the 'leading edge' on cutting tools, the hardness is derived from the very strong rigid three–dimensional carbon–carbon bond network.

o Diamond also has a very high melting point because of this very strong giant covalent lattice in which every carbon atom is strongly bonded to four other carbon atoms (see diagram above on right).

o The strong bond network in diamond (and graphite and silica) prevents these materials from dissolving in any conventional solvent.

o Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances is given in a section of the Energetics Notes.

o Pure elemental silicon (not the oxide) has the same molecular structure as diamond and similar properties, though not as strong or high melting.

SILICON DIOXIDE (SILICA) (SiO2)n where n is an extremely

large number of silicon and oxygen atoms!o Many naturally occurring minerals are based on –

O–X–O– linked 3D structures where X is often silicon (Si) and aluminium (Al), three of the most abundant elements in the earth's crust.

o Silicon dioxide ('silica') is found as quartz in granite (igneous rock) and is the main component in sandstone – which is a sedimentary rock formed the compressed erosion products of igneous rocks.

Looking at the diagram on the right of silica, each silicon atom (black blobs) forms four strong covalent bonds with the linking oxygen atoms (yellow blobs).

Again like diamond, theoretically all the atoms in a silica crystal are linked together by a strong 3D covalent bond network.

o Silica (SiO2) is a very hard substance with a very high melting point and won't dissolve in any solvent.

o There are no free electrons so silicon dioxide doesn't conduct electricity.

o Many more minerals that are hard wearing, rare and attractive when polished, hold great value as gemstones, but sand is also mainly silica, but not quite as valuable!

GRAPHITE multilayers of Cn sheets where n is an extremely large number of carbon atoms!

o Carbon also occurs in the form of graphite. The carbon atoms form joined hexagonal rings

forming layers 1 atom thick in graphite. A crystal of graphite contains millions of layers

of these sheets of carbon atoms. Although graphite is almost black and opaque

(unlike diamond), it does look a bit shiny and smooth.

o There are three strong covalent bonds per carbon atom in graphite (3 C–C bonds in a planar arrangement from 3 of its 4 outer electrons). So three of the electrons are tightly held in three directed covalent bonds, BUT, the fourth outer electron is 'delocalised' or shared between the carbon atoms to form the equivalent of a 4th bond per carbon atom AND is free to move around - hence graphite's ability to conduct electricity.

This situation requires advanced level concepts to fully explain the structure of graphite, and this bonding situation also occurs in fullerenes described below, and in aromatic compounds you deal with only at advanced level.

o The layers are only held together by weak intermolecular forces shown by the dotted lines NOT by strong covalent bonds, so graphite, for a giant covalent structure, is unusually weak physically.

o Like diamond and silica (above) the large molecules of the layer ensure graphite has typically very high melting point because of the strong 2D bonding

GRAPHITE

network (note: NOT 3D network).o Graphite will not dissolve in solvents because of

the strong bonding in the layers.o BUT there are two crucial differences compared to

diamond ... Electrons, from the 'shared bond', can

move freely through each layer, so graphite is a conductor like a metal(diamond is an electrical insulator and a poor heat conductor). Graphite is used in electrical contacts e.g. electrodes in electrolysis.

The weak forces enable the layers to slip over each other so where as diamond is hard material graphite is a 'soft' crystal, it feels slippery.

This enables graphite to be used as a lubricant.

Carbon in the form of graphite is the only non–metal that is a significant electrical conductor.

Graphite is used in pencils (often wrongly called lead pencils!) because the weak structure allows the layer to slide off onto paper when pressure is applied on rubbing the pencil over paper.

o These two different characteristics of graphite described above are put to a common use with the electrical contacts in electric motors and dynamos.

These contacts (called brushes) are made of graphite sprung onto the spinning brass contacts of the armature.

The graphite brushes provide good electrical contact and are self–lubricating as the carbon layers slide over each other.

A 3rd form of carbon (another allotrope of carbon) are fullerenes or 'bucky balls'! They consists of hexagonal rings like graphite and alternating pentagonal rings to allow curvature of the surface, in fact curved sufficiently to form 'football' or 'rugby ball' shapes..

Buckminster Fullerene C60 is shown on the right and the bonds form a pattern like a soccer ball.

o Others are oval shaped like a rugby ball. It is a black solid insoluble in water.

They are NOT considered giant covalent structures and are classed as simple molecules. They do dissolve in organic solvents giving coloured solutions (e.g. deep red in petrol hydrocarbons, and although solid, their melting points are not that high.

They are mentioned here to illustrate the different forms of carbon AND they can be made into continuous tubes to form very strong fibres of 'pipe like' molecules called 'nanotubes'. These 'molecular size' particles behave quite differently to a bulk carbon material like graphite.

Uses of Nanotubes – carbon nanotechnology – examples of nanochemistry

FULLERENES

o They can be used as semiconductors in electrical circuits.

o They act as a component of industrial catalysts for certain reactions whose economic efficiency is of great importance (time = money in business!).

The catalyst can be attached to the nanotubes which have a huge surface are per mass of catalyst 'bed'.

They large surface combined with the catalyst ensure two rates of reaction factors work in harmony to increase the speed of an industrial reaction so making the process more efficient and more economic.

o Nanotube fibres are very strong and so they are used in 'composite materials' e.g. reinforcing graphite in carbon fibre tennis rackets.

o Nanotubes can 'cage' other molecules and can be used as a means of delivering drugs in controlled way to the body because the thin carbon nanotubes can penetrate cell walls.

METALLIC BONDING – structure and properties of metals

 Explaining the physical properties of metals

All metals are lustrous and, compared to non-metals, most metals are quite dense, hard (tough, high tensile strength), with high melting/boiling points, though there notable exceptions e.g.

o mercury is a liquid at room temperature, group 1 alkali metals like sodium and potassium are less dense than water ('float') and have low melting points <100oC).

The strong bonding generally results in dense, strong materials with high melting and boiling points.

o Usually a relatively large amount of energy is needed to melt or boil metals.o The stronger the attraction between the atoms/ions in the giant metallic lattice, the

more energy is needed to weaken the force between them sufficiently to break the giant lattice down in melting and completely to boil the metal.

o The strong bonding in metals gives them a high tensile strength, so alloys like steel are used in building construction, car bodies etc.

Metals are good conductors of electricityo Why are metals good conductors of electricity? Metals are good at conducting

electricity because these 'free' electrons carry the charge of an electric current when a potential difference (voltage!) is applied across a piece of metal

e.g. copper wire is excellent to use as an electrical conductor in household wiring or any electrical appliances.

Metals are also good conductors of heat.o Why are metals good conductors of heat? The fact that metals are good at

conducting heat is also due to the free moving electrons.o Non–metallic solids conduct heat energy by hotter more strongly vibrating atoms,

knocking against cooler less strongly vibrating atoms to pass the particle kinetic energy on.

o BUT in metals, as well as this effect, the 'hot' high kinetic energy electrons move around freely to transfer the particle kinetic energy more efficiently to 'cooler' atoms. This is a faster process than the transferring heat by the kinetic energy of atom vibration.

So, where a material needs to be a good heat conductor, metals quite naturally are used to make everything from radiators, cooking pans etc.

Its also hand that they are both strong and high melting when used as a saucepan!

Typical metals also have a silvery surface (lustrous) but remember this may be easily tarnished by corrosive oxidation in air and water.

o Although many metals will corrode (oxidise) in the presence of air (oxygen) and water, the strong bonding prevents them dissolving in water or any other laboratory solvent. When metals like sodium 'dissolve in water, they do so via a chemical reaction forming a soluble compound (sodium hydroxide), and do NOT give a solution of sodium metal.

Unlike ionic solids, metals are very malleable - easy to bend or hammer into shapeo Why are metals very malleable and easily bent or pressed shaped? Metals are be

readily bent, pressed or hammered into shape because the strong bonding is retained even when the metal is stressed (at least upto a point!)

The layers of atoms can slide over each other without fracturing the structure.

The reason for this is the mobility of the electrons involved in the metallic bonding.

When planes of metal atoms are 'bent' or slide the electrons can run in between the atoms and maintain a strong bonding situation. This can't happen in ionic solids which tend to be brittle and the layers of immobile ions fracture easily.

Potential problems with metal structures?o Although the metals used in construction are strong, in some situations they may

become dangerously weak e.g. If iron or steel becomes badly corroded, there is no strength in rust!, and,

the thicker the rust layer, the thinner and weaker the supporting iron or steel layer, hence the possibility of structural failure. Therefore, most iron and steel structures exposed to the outside weather are maintained with a good coating of paint.

Also, if metal structures e.g. in aircraft or bridges, are continually strained under stress, the crystal structure of the metal can change so it becomes brittle. This effect is called metal fatigue (stress fractures) and may lead to a very dangerous situation of mechanical failure of the structure.

So it is important develop alloys which are well designed, well tested and will last the expected lifetime of the structure whether it be part of an aircraft (eg titanium aircraft frame) or a part of a bridge (eg steel suspension cables).

 Note on Alloy Structure  via a very simplified diagram

An alloy is a mixture of a metal with other elements (metals or non-metals). Metals can be mixed together to make alloys to improve the metal's properties to better suit a particular

purpose. An alloy mixture often has superior desired properties compared to the pure metal or metals i.e. the alloy has its own unique properties and a more useful metal. 

1. Shows the regular arrangement of the atoms in a pure metal crystal and the white spaces show where the free electrons are (yellow circles actually positive metal ions).

2. Shows what happens when the metal is stressed by a strong force. The layers of atoms can slide over each other and the bonding is maintained as the mobile electrons keep in contact with ions of the giant lattice, so the metal object remains intact BUT the metal is physically a different shape.

3. Shows an alloy mixture. Alloys are NOT compounds but a physical mixing of a metal plus at least one other material (shown by red circle), it can be another metal e.g. nickel or manganese added to iron in steel, or a non–metal e.g. carbon, and it can be bigger or smaller than the iron atoms. Many alloys are produced like this to give a stronger metal. The presence of the other atoms (smaller or bigger) disrupts the symmetry of the layers and reduces the 'slip ability' of one layer next to another. The result is a stronger harder less malleable metal.

4. The main point about using alloys is that you can make up, and try out, all sorts of different compositions until you find the one that best suits the required purpose in terms of tensile/compression strength, malleability, electrical conductivity or corrosion resistance etc.

o The are hundreds of alloys of steel made by alloying iron with other metals to increase the strength or anti-corrosion properties of the metal.

Steel is used in building and bridge construction, car bodies, railway lines and countless other objects that need to have a high tensile strength.

o Pure metals can be either too soft (e.g. like copper or tin) or too brittle (e.g. like zinc) to be used directly and are therefore often alloyed to make superior metals like brass or bronze.

o The properties of metals are readily matched to a particular use e.g. Aluminium alloys are strong and light (relatively low density for a

metal), they do not corrode easily and so are used in aircraft construction, greenhouse frames and not as expensive as titanium alloys.

Cooking pans made of stainless steel are good conductors of heat, strong with good anti-corrosion properties and steel has a high melting point.

Copper is malleable and ductile, easily drawn out into wire, an excellent conductor of electricity, and so is widely used in electrical circuitry.

o Steel alloys of varying strength and anti-corrosion properties are used in thousands of products and constructions e.g. reinforcing rods in concrete buildings, bridge girders, car engines, domestic appliances from washing machines to electric kettles, saucepans, tools like chisels, ship hulls and superstructure, very hard drill bits,

The crystal structure and properties of Ionic Compounds