module 5 chemical bonds.doc

7/30/2019 Module 5 Chemical Bonds.doc

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MODULE 5


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Why do atoms bond together? - 'electron glue'!

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).

(a) IONIC BONDING - By one atom transferring electrons to another atom to form

oppositely charged particles called ions which attract each other - the ionic bond.

An ion is an atom or group of atoms carrying an overall positive or negative

charge

o E.g. Na+, Cl-, [Cu(H2O)]2+, SO4

2- etc.

If a particle, as in a neutral atom, has equal numbers of protons (+) and electrons

(-) the particle charge is zero i.e. no overall electric charge.

The proton/atomic number in an atom does not change BUT the number of

associated electrons can!

If negative electrons are lost the excess charge from the protons produces an

overall positive ion.

If negative electrons are gained there is an excess of negative charge, so a negative

ion is formed.


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The charge on the ion is numerically related to the number of electrons transferred i.e.

electrons lost or gained.

For any atom or group of atoms, for every electron gained you get a one unit

increase in negative charge on the ion, for every electron lost you get a one unit

increase in the positive charge on the ion.

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

The atom gaining electrons forms a negative ion (anion) and is usually a non-

metallic element. The ionic bond then consists of the attractive force between the

positive and negative ions in the structure.

The ionic bonding forces act in all directions around a particular ion, it is not

directional, as in the case of covalent bonding.

(b) COVALENT BONDING - sharing electrons to form molecules with covalent bonds,

the bond is usually formed between two non-metallic elements in a molecule. The two

positive nuclei (due to the positive protons in them) of both atoms are mutually

attracted to the shared negative electrons between them - the covalent bond. They share

the electrons in a way that gives a stable Noble Gas electron arrangement.

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 molecules

have a particular shape.

(c) METALLIC BONDING isn't quite like ionic or covalent bonding, the metal atoms

form positive ions, but no negative ion is formed from the same metal atoms, but the

positive metal ions/atoms are attracted together by the free moving negative electronsbetween them.

NOBLE GASES are very reluctant to share, gain or lose electrons to form a chemical

bond. They are already electronically very stable. For most other elements the types of

bonding and the resulting properties of the elements or compounds are described in

detail in Parts 2 to 5. In all the electronic diagrams ONLY the outer electrons are

shown.

IONIC BONDING - compounds and properties

Examples of ionic compounds*physical properties of ionic compounds

Ionic Bonding - electron transfer

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 thenumber of positive protons equals the number of negative electrons .


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Ions are atoms, or groups of atoms, which have lost or gained electrons to have a net

electrical charge overall.

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).

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.

The examples below combining a metal from Groups 1 (Alkali Metals), 2 or 3, with a non-

metal from Group 6 or Group 7 (The Halogens). The electron structures are shown in () or [].

Only the outer electrons of the original atoms, and where they end up in the ions, are shown

in the dot and cross (ox) diagrams

Ionic bonding is not directional like covalent bonding, in the sense that the force of attractionbetween the positive ions and the negative ions act in every direction around the ions.

Example 1: A Group 1 metal + a Group 7 non-metal e.g. sodium + chlorine ==> sodium

chloride NaCl or ionic formula Na+Cl- In terms of electron arrangement, 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, sodium

becomes like neon and chlorine like argon.

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]-

ONE combines with ONE to form

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

similar.

Note:

The charge on the sodium ion Na+ is +1 units (shown as just +) because there is one more

positive proton than there are negative electrons in the sodium ion.

The charge on the chloride ion Cl- is -1 unit (shown as just -) because there is one more

negative electron than there are positive protons in the chloride ion.


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Example 2: A Group 2 metal + a Group 7 non-metal e.g. magnesium + chlorine ==>

magnesium chloride MgCl2 or ionic formula Mg2+(Cl-)2 In terms of electron arrangement,

the magnesium donates its two outer electrons to two chlorine atoms forming a double

positive magnesium ion and two single negative chloride ions. 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

ONE combines with TWO to form see *

* 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 BeF 2, magnesium

bromide MgBr2, calcium chloride CaCl2 or barium iodide BaI2 etc. will all be electronically

similar.

Ca is 2.8.8.2, F is 2.7 rest of dot and cross diagrams are up to you.

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, thealuminium 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. 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

ONE combines with THREE to form

Solid aluminium chloride/bromide/iodide have similar formula but are covalent when

vapourised into Al2X6 dimer molecules.

Example 4: A Group 1 metal + a Group 6 non-metal e.g. sodium/potassium + oxygen ==>

sodium/potassium oxide Na2O/K2O or ionic formula (Na+)2O2-/(K+)2O2- In terms of

electron arrangement, the two sodium/potassium atoms donate their outer electron to one


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oxygen atom. This results in two single positive potassium ions to one double negative oxide

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

Valencies, K 1, oxygen 2. Lithium oxide, Li2O, sodium oxide Na2O, sodium sulphide Na2S

and potassium K2S etc. will be similar.

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-

TWO combine with ONE to form

or

+ ==>

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-

TWO combine with ONE to form

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 dots and crosses

diagrams are up to you.


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Example 5: A Group 2 metal + a Group 6 non-metal e.g. magnesium/calcium + oxygen

==> magnesium/calcium oxide MgO/CaO or ionic formula Mg2+O2-/Ca2+O2- In terms of

electron arrangement, 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.

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


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] + [2,6] ==>

[2,8,8]2+ [2,8]2-

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-


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).

Formagnesium sulphide: Mg (2.8.2) + S (2.8.6) ==> Mg2+ (2.8) S2- (2.8.8)

Forcalcium sulphide: Ca (2.8.8.2) + S (2.8.6) ==> Ca2+ (2.8.8) S2- (2.8.8)

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).


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Example 6: A Group 3 metal + a Group 6 non-metal e.g. aluminium + oxygen ==>

aluminium oxide Al2O3 or ionic formula (Al3+)2(O

2-)3 In terms of electron arrangement, 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. 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

TWO combine with THREE to form

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.

The properties of Ionic Compounds

The diagram on the right is typical of the

giant ionic crystal structure of ionic

compounds like sodium chloride and

magnesium oxide.

The alternate positive and negative ions in

an ionic solid are arranged in an orderlyway in a giant ionic lattice structure shown

on the left.

The ionic bond is the strong electrical attraction between the 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 has

high melting and boiling points, so they are not very volatile!


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A relatively large amount of energy is needed to melt or boil ionic

compounds.

The bigger the charges on the ions the stronger the bonding attraction

e.g. magnesium oxide Mg2+O2- has a higher melting point than sodium

chloride Na+

Cl-

.

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. They are NOT malleable like metals

(see below).

Many ionic compounds are soluble in water, but not all, so don't make

assumptions. Salts can dissolve in water because the ions can separate

and become surrounded by water molecules which weakly bond to the

ions. This reduces the attractive forces between the ions, preventing thecrystal structure to exist. 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. However, if the ionic

compound is melted or dissolved in water, the liquid will now conduct

electricity, as the ion particles are now free.

3. Covalent Bonding - electron sharing in big or small molecules!

Covalent bonds are formed by atoms sharing electrons to form molecules. 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.

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

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 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 molecules have a


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particular shape. In the case of ionic or metallic bonding, the electrical attractive forces act

in all directions around the particles involved.

The bonding in Small Covalent Molecules

The simplest molecules are formed from two atoms and examples of their formation are

shown below. The electrons are shown as dots and crosses to indicate which atom the

electrons come from, though all electrons are the same. The diagrams may only show

the outer electron arrangements for atoms that use two or more electron shells. The

electron structures are given in (). Examples of simple covalent molecules are

Example 1: two hydrogen atoms (1) form the molecule of the element hydrogen H2

and combine to form where both atoms have a pseudo helium

structure of 2 outer electrons around each atom's nucleus. Any covalent bond is formed from

the mutual attraction of two positive nuclei and negative electrons between them. The nuclei

would be a tiny dot in the middle of where the H symbols are drawn! H valency is 1.

Example 2: two chlorine atoms (2.8.7) form the molecule of the element chlorine Cl2

and combine to form where both atoms have a

pseudo argon structure of 8 outer electrons around each atom. All the other halogens would

be similar e.g. F2, Br2 and I2 etc. Valency of halogens like chlorine is 1 here.

Example 3: one atom of hydrogen (1) combines with one atom of chlorine (2.8.7) to form the

molecule of the compound hydrogen chloride HCl

and combine to form where hydrogen is electronically like

helium and chlorine like argon. All the other hydrogen halides will be similar e.g. hydrogen

fluoride HF, hydrogen bromide HBr and hydrogen iodide HI.


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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 ofhydrogen ions (H+) and chloride ions (Cl-). The solution then conductselectricity and passage of a d.c. current causes electrolysis to take place forming hydrogen

and chlorine.

Example 4: two atoms of hydrogen (1) combine with one atom of oxygen (2.6) to form the molecule

of the compound water H2O

and and combine to form so that the hydrogen atoms are

electronically like helium and the oxygen atom becomes like neon. The molecule can be shown as

with two hydrogen - oxygen single covalent bonds (AS note: called a V or bent shape, the

H-O-H bond angle is 105o). 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.

Example 5: three atoms of hydrogen (1) combine with one atom of nitrogen (2.5) to form themolecule of the compound ammonia NH3

three of and one combine to form so that the hydrogen

atoms are electronically like helium and the nitrogen atom becomes like neon. The molecule

can be shown

as with three nitrogen - hydrogen single covalent bonds (AS note: called a trigonal pyramid

shape, the H-N-H bond angle is 107o). 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.

Example 6: four atoms of hydrogen (1) combine with one atom of carbon (2.4) to form themolecule of the compound methane CH4


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four of and one of combine to form so that the hydrogen

atoms are electronically like helium and the nitrogen atom becomes like neon. The molecule

can beshown as 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).

Example 7: Two atoms of oxygen (2.6) combine to form the molecules of

the element oxygen O2. The molecule has one O=O double covalent bond .

Oxygen valency 2.

Example 8: One atom of carbon (2.4) combines with two

atoms of oxygen (2.6) to form the compound carbon dioxide CO2. The molecule can be

shown as 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.


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Example 9: Two atoms of carbon (2.4) combine with four atoms of

hydrogen (1) to form ethene C2H4. The molecule can be shown as with one

carbon = carbon double bond and four carbon - hydrogen single covalent bonds (called a planar

shape, its completely flat!, the H-C=C and H-C-H bond angles are 120o). The valency of carbon is still

4.

Examples 10-13: The scribbles below illustrate some more complex examples. Can you

deduce them for yourself? Ex. 10nitrogen N2; Ex. 11ethane C2H6; Ex. 12chloromethane

CH3Cl and Ex. 13 methanol CH3OH. Electronic origin of the diagrams showing the outer

electrons of N, C, Cl and O: N at. no. 7 (2.5), H at. no. (1), C at. no. 6 (2.4), Cl at. no. 17

(2.8.7) and O at. no. 8 (2.6) plus a variety of crosses and blobs! The valencies or combining

power in these examples are N 3, H 1, C 4, Cl 1 and O 2. From these you can work out others

e.g. Ex. 12 can be used to derive the ox diagram for tetra-chloromethane CCl4.

AS advanced level notes on shapes and bond angles:

o Ex. 11 Ethane has a linked double tetrahedral shape, all H-C-H and H-C-C

bond angles are 109o

o Ex. 12 chloromethane has tetrahedral shape with H-C-H and H-C-Cl bond

angles of approximately 109o

o Ex. 13 methanol, the four bonds around the central carbon are tetrahedrally

arranged with a H 'wiggle' on the oxygen. All the H-C-H, H-C-O and C-O-H

bond angles are approximately 109o

o The blue icon e.g. below, represents an octahedral shape (e.g. SF6, complex

transition metal ions like [Cu(H2O)6]2+ and the bond angles are either 90o or

180o

o Simple molecules with a triple bond are often linear e.g. H-C C-H ethyne

orH-C N hydrogen cyanide (methanenitrile)


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Typical properties of simple covalent substances - small molecules!

Covalent substances tend to be liquids or gases at room temperature because the forces

between their particles (the molecules) are weak. These intermolecular forces must not be

confused with intramolecular forces. Inter- means between; intra- means within.

Intermolecular forces are forces between individual molecules (such as van der Waal

forces). They are weak forces.

Intramolecular forces are forces within the molecules (covalent bonds). They hold the

atoms together in the molecule. These are strong forces.

-

+ +

-

Intra-molecular force +

Intermolecular force (between one water molecule and another)

-

+ +

Intermolecular forces

We will discuss two types of intermolecular forces:

van der Waals forces

hydrogen bonds

Van der Waals forces

All covalent molecules, whether polar or non polar, develop temporary or instantaneous

dipoles. This results from uneven movement of all the electrons within the molecules. Van

der waals forces are the weak attraction between oppositely charged ends of molecules with

temporary dipoles.

H

s

s

s

s

H

O

H H

O

H

H

O


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+ - + - + -

a molecule a molecule ..but can develop one

with a temporary that does

dipole have a dipole

Figure 4.0 Van der Waals forces. When molecule 2 develops a temporary dipole then it will be attracted

to molecule 1.

HYDROGEN BONDS

The hydrogen bond is the weak attraction between the slightly positive hydrogen atom in one

polar molecule and the slightly electronegative atom in another polar molecule of the same

type or of a different type.

Structure and Properties of simple covalent substances

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 I-I covalent bond is

strong enough to withstand the heating and the purple vapour still consists of

the same I2 molecules as the dark coloured solid is made up of.

So why the ease of vaporisation on heating?

o The electrical attractive forces between individual molecules are weak, so

the bulk material is not very strong physically and there are also

consequences for the melting and boiling points.

These weak electrical attractions are known as intermolecular forces and are

readily weakened further on heating. 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 On heating the inter-molecular forces are easily overcome with the increased

kinetic energy of the particles giving the material a low melting orboiling

point and a relatively small amount of energy is needed to effect these state

changes.

o This contrasts with the high melting points of giant covalent structures with

their strong 3D network.

o Note: 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 aresometimes referred to as intramolecular forces (i.e. internal to the molecule).

12

12


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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 This again contrasts with giant covalent structures where the strong bond

network stops solvent molecules interacting with the particles making up the

material.

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!

Large Covalent Molecules and their Properties

(Macromolecules - giant covalent networks and polymers)

Because covalent bonds act in a particular direction i.e. along the 'line' between the two

nuclei of the atoms bonded together in an individual bond, strong structures can be

formed, especially if the covalent bonds are arranged in a strong three dimensional

giant covalent lattice.

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

DIAGRAMS


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It is possible formany atoms to link up to form a giant

covalent structure or lattice. The atoms are usually non-

metals.

This produces a very strong 3-dimensional covalent

bond network or lattice.

This gives them significantly different properties from

the small simple covalent molecules mentioned above.

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. 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 smallS8 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 and has a very high melting and boiling points

because of the strong covalent bond network (3D or 2D inthe case of graphite below).

A relatively large amount of energy is needed to melt or

boil giant covalent structures.

They are usually poor conductors of electricity because

the electrons are not usually free to move as they can in

metallic structures.

Also because of the strength of the bonding in all

directions in the structure, they are often very hard,

strong and will not dissolve in solvents like water. The

DIAMOND

SILICA

silicon dioxide


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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.

The hardness of diamond enables it to be used as the

'leading edge' on cutting tools.

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 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.

o Many some minerals that are hard wearing, rare

and attractive when polished hold great value as

gemstones.

Carbon also occurs in the form of graphite. The carbon

atoms form joined hexagonal rings forming layers 1 atom

thick.

There are three strong covalent bonds per carbon (3

C-C bonds in a planar arrangement from 3 of its 4 outerelectrons), BUT, the fourth outer electron is 'delocalised'

or shared between the carbon atoms to form the

equivalent of a 4th bond per carbon atom (this situation

requires advanced level concepts to fully explain, and this

bonding situation also occurs in fullerenes described

below, and in aromatic compounds you deal with at

advanced level).

The layers are only held together by weak

intermolecular forces shown by the dotted lines NOT by

strong covalent bonds.

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 network(note:

NOT 3D network)..

Graphite will not dissolve in solvents because of the

strong bonding

BUT there are two crucial differences compared to

diamond ...

o Electrons, from the 'shared bond', can move

GRAPHITE


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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.

o 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.

Graphite is used as a lubricant.

These two different characteristics 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 asthe carbon layers slide over each other.

A 3rd form of carbon are fullerenes or 'bucky balls'! It

consists of hexagonal rings like graphite and alternating

pentagonal rings to allow curvature of the surface.

Buckminster Fullerene C60 is shown and the bonds form

a pattern like a soccer ball. 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:

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

FULLERENES


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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

the industrial reaction.

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.

BONDING IN METALS

METALLIC BONDING - structure and properties of metals

The crystal lattice of metals consists of ions NOTatomssurrounded by a 'sea of electrons' forming

another type ofgiant lattice.

The outer electrons(-) from the original metal atoms are free to move

around between the positive metal ions formed (+).

These free or 'delocalised' electrons are the 'electronic glue' holding the

particles together.

There is a strong electrical force of attractionbetween these mobile

electrons (-) and the 'immobile' positive metal ions (+) and this is themetallic bond.

Metallic bonding is not directional like covalent bonding, it is like ionic

bonding in the sense that the force of attraction between the positive metal

ions and the mobile electrons acts in every direction about the fixed

(immobile) metal ions.

Explaining the physical properties of metals

This strong bonding generally results in dense, strong materials with high melting

and boiling points. Usually a relatively large amount of energy is needed to melt or boil metals. .


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Metals are good conductors of electricitybecause these 'free' electrons carry the

charge of an electric current when a potential difference (voltage!) is applied across a

piece of metal.

Metals are also good conductors of heat. This is also due to the free moving

electrons. Non-metallic solids conduct heat energy by hotter more strongly vibratingatoms, knocking against cooler less strongly vibrating atoms to pass the particle

kinetic energy on. 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.

Typical metals also have a silvery surface but remember this may be easily tarnished

by corrosive oxidation in air and water.

Unlike ionic solids, metals are very malleable, they can be readily bent, pressed or

hammered into shape. The layers of atoms can slide over each other without

fracturing the structure (see below). The reason for this is the mobility of the

electrons. 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.

Note on Alloy Structure

1. Shows the regular arrangement of the atoms in a 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 atoms, so the metal remains intact BUT

a different shape.

3. Shows an alloy mixture. It is NOT a compound but a physical mixing of a

metal plus at least one other material (shown by red circle, it can be another

metal e.g. Ni, a non-metal e.g. C or a compound of carbon or manganese, and

it can be bigger or smaller than iron atoms). Many alloys are produced 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.


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POLAR AND NON-POLAR COVALENT COMPOUNDS

POLAR MOLECULES

A covalent bond in which the electron pair is shared unequally has partial ionic character and

is called a polar covalent bond. ( A polar covalent is a bond between two atoms that

have partial electric charges arising from their difference in electro-negativity. The

partial charges give rise to an electric dipole moment.)The extent to which an atom has a

greater or lesser share of electrons in a bond is determined by the difference in electro-

negativities of the two bonds. The electro-negativity of an element is its electron-pulling

power when it is part of a compound. An atom with a high electro-negativity has a strong

pulling power on electrons particularly for electron pair it shares with its neighbour. The

outcome of the tug-of-war: the more electronegative atom has a greater share of electron pair

of the covalent bond.

An O-H bond is polar because oxygen is more electronegative than hydrogen and gains a

greater share in the bonding electron pair. Its greater share of electrons means that oxygen

has a partial negative charge, which we denote -. Because the electron pair has been

pulled away from hydrogen atom, that atom has a partial positive charge, denoted +. We

show the partial charges on the atoms by writing +H-O - . A polar molecule is a molecule

with a nonzero dipole moment. All diatomic molecules composed of atoms of different

elements are slightly polar.

Examples of polar molecules

HF

HCl

HBr

HI - +

CO

ClF

H2O

NH3


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NON-POLAR MOLECULES

A non-polar molecule is molecule that has zero electric dipole moment. All homo-nuclear

(same) diatomic molecules, such as Cl2 and H2, are non-polar because there are no partial

charges on their atoms.

Examples of Non- polar Molecules

Cl2

H2

F2

N3

O2

module 5 chemical bonds.doc

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