DOC BROWN'S Science-CHEMISTRY HOMEPAGE KS3 SCIENCE QUIZZES and WORKSHEETS (~US grades 6-8)
GCSE SCIENCE help links GCSE ADDITIONAL SCIENCE help links
KS3 BIOLOGY Quizzes KS3 CHEMISTRY Quizzes & Worksheets KS3 PHYSICS Quizzes
KS4 Science GCSE/IGCSE CHEMISTRY NOTES (~US grades 8-10) KS4 Science GCSE/IGCSE CHEMISTRY QUIZZES and WORKSHEETS (~US grades 8-10) ADVANCED LEVEL CHEMISTRY QUIZZES and WORKSHEETS (~US grades 11-12)
Custom Search

(c) doc b(c) doc b

CHEMICAL BONDING Part 3 Covalent Bonding in ..

small simple molecules, simple molecular substances

Doc Brown's Science–Chemistry Chemical Bonding GCSE/IGCSE/O/AS/A2 Level  Revision Notes

What is the bonding in simple molecules? COVALENT BONDING DIAGRAMS and PROPERTIES OF SIMPLE SMALL COVALENT MOLECULES The in bonding small covalent molecules is described and how covalent bonds are formed by sharing electrons and the formation of small simple molecules and their physical properties of simple molecular substances are described, discussed and explained. Examples include hydrogen, chlorine, hydrogen chloride, water, ammonia, methane, oxygen, carbon dioxide, ethene, nitrogen, ethane, chloromethane, methanol. Also, a section on how to work out the molecular formula of a covalent compound from the valencies of the constituent atoms.  These notes on covalent bonding in molecules of simple molecular substances and explaining their properties are designed to meet the highest standards of knowledge and understanding required for students/pupils doing GCSE chemistry, IGCSE chemistry, O Level chemistry, KS4 science courses and a basic primer for AS/A Level chemistry courses.

Part 1 Introduction – why do atoms bond together? (I suggest you read 1st)

Part 2 Ionic Bonding – compounds and properties

Part 3 Covalent Bonding – small simple molecules and properties (this page)

See also  Appendix 1. More on intermolecular forces – intermolecular bonding

and Appendix 2. How to work out a covalent compound formula

Part 4 Covalent Bonding – macromolecules and giant covalent structures

Part 5 Metallic Bonding – structure and properties of metals

Part 6 More advanced concepts for advanced level chemistry (in preparation, BUT a lot on intermolecular forces in Equilibria Part 8)



Part 3. COVALENT BONDING small simple molecules and properties

simple small molecule bonding e.g. water * physical properties of small molecules

inter/intra (internal)–molecular forces

Examples of covalent elements/compound examples described: hydrogen H2, chlorine Cl2, hydrogen chloride HCl, water H2O, ammonia NH3, methane CH4, oxygen O2, carbon dioxide CO2, ethene C2H4, nitrogen N2, ethane C2H6, chloromethane CH3Cl, methanol CH3OH

Advanced level chemistry notes on the shapes and bond angles of molecules and ions

See also  Appendix 1. More on intermolecular forces – intermolecular bonding

and Appendix 2. How to work out a covalent compound formula

top


(c) doc b3. Covalent Bonding – electron sharing in big or small molecules!

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!

(c) doc b (c) doc b (c) doc b 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.

(c) doc b

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

Pd metals Part of the modern Periodic Table

Pd = period, Gp = group

metals => non–metals
Gp1 Gp2 Gp3 Gp4 Gp5 Gp6 Gp7 Gp0
1

1H  Note that H does not readily fit into any group

2He
2 3Li 4Be atomic number Chemical Symbol eg 4Be 5B 6C 7N 8O 9F 10Ne
3 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar
4 19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
6 55Cs 56Ba Transition Metals 81Tl 82Pb 83Bi 84Po 85At 86Rn
Gp 1 Alkali Metals  Gp 2 Alkaline Earth Metals  Gp 7 Halogens  Gp 0 Noble Gases

Chemical 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 formation of a covalent compounds like hydrogen chloride HCl, sulfur dioxide SO2, and element molecules like hydrogen H2 or oxygen O2.

 

Most covalent molecules you will come across are formed by combinations of atoms of non–metallic elements on the right–hand side of the Periodic Table eg from Group 4 carbon and silicon, from Group 5 nitrogen and phosphorus and from the Group 7 Halogens – fluorine, chlorine, bromine and iodine. Hydrogen also forms predominantly covalent compounds and so does the Noble Gas xenon in Group 0. Don't forget, even non–metal elements can form molecules eg hydrogen H2 and phosphorus P4.

Brief summary of the Periodic Table including electronic structure and formula patterns

top


(c) doc bThe 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 parentheses ().

Examples of simple covalent molecules are …


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.

(c) doc b and (c) doc b combine to form (c) doc b 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.

(c) doc b and (c) doc b combine to form (c) doc b where both atoms have a pseudo argon structure of 8 outer electrons around each atom.

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)..

(c) doc bElectronically, 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.

top


Example 3Covalent 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.

(c) doc b and (c) doc b combine to form (c) doc b 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.

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.

Reminder: How to work out formula of covalent compounds without going through some demanding electronic thinking is described on the "Elements, Compounds and Mixtures" page and it is followed by a section on naming compounds.

top


Example 4Covalent 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.

(c) doc b and (c) doc band (c) doc b combine to form (c) doc b 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 (c) doc b(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.


Example 5Covalent 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 (c) doc b and one (c) doc b combine to form (c) doc b 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 (c) doc b (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.

top


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 (c) doc b and one of (c) doc b combine to form (c) doc b 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.

(c) doc bNote 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 (c) doc b (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.


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

(c) doc b 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 (c) doc b(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.

top


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

 (c) doc b 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 molecule can be shown as (c) doc b (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

(c) doc b 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 (c) doc b (displayed formula) 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.

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.

top


Examples 10–13: The scribbles below illustrate some more complex examples with their molecular formula.

Only the outer electrons for carbon, nitrogen, oxygen and chlorine are shown.

Can you deduce these electronic diagrams of covalently bonded molecules for yourself?

Ex. 10 nitrogen N2; Ex. 11 ethane C2H6; Ex. 12 chloromethane 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 theses 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 tetrachloromethane CCl4.

(c) doc b

Two other electronic diagrams for Ex. 11 ethane and

and its displayed formula alkanes structure and naming (c) doc b

Alkanes are relatively small molecules in which all the chemical bonds are covalent bonds. All the bonds in alkane molecules are single bonds i.e. C–C carbon – carbon or C–H carbon – hydrogen. Each carbon atom forms four single covalent bonds and hydrogen atoms form one single covalent bond. All single bonds are formed by sharing a pair of electrons e.g. one from each of a carbon atom and a hydrogen atom, or two carbon atoms contributing (sharing) an electron each to the covalent bond.

  • AS advanced level notes on shapes and bond angles:

    • Ex. 11 Ethane has a linked double tetrahedral shape, all H–C–H and H–C–C bond angles are 109o

    • Ex. 12 chloromethane has tetrahedral shape with  H–C–H and H–C–Cl bond angles of approximately 109o

    • 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

    • 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

    • Simple molecules with a triple bond are often linear e.g. H–C(c) doc bC–H ethyne or H–C(c) doc bN hydrogen cyanide (methanenitrile)

    • The theory of shapes and bond angles with more examples and diagrams is on another page for AS–A2 students and with an extra section on bond angles in organic molecules.

  • on another web page is  how to work out a covalent formula given the element valencies (combining power)

top


(c) doc bTypical properties of simple molecular substances

Composed of relatively small covalently bonded molecules
(c) doc b (c) doc b (c) doc b

  • Why do simple covalent molecules typically have low melting and low boiling points?

    • Typical examples are water (ice & liquid), petrol, butter etc.

    • 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'.

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

    • 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.

    • 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.

    • 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.

    • 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.

    • 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?

    • 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.

    • 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 (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.

    • 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.

    • 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.

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

    • 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.

    • 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).

    • 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!

  • More advanced notes on intermolecular forces

  • More advanced level chemistry notes on the shapes and bond angles of molecules and ions


Appendix 1A little more on intermolecular forces – intermolecular bonding

using water as an example

Between all particles, but with particular reference to covalently bonded molecules, there always exists some very weak electrical attractive forces known as intermolecular forces or intermolecular bonding.

These constantly acting attractive forces or intermolecular bonds are very much weaker than covalent or ionic chemical bonds (approximately 1/30 to 1/20th in comparative attractive force).

For example, although the oxygen and hydrogen atoms are very strongly bonded in water to make a VERY stable molecule, BUT this does NOT account for the existence of liquid water and ice!

It is the weak intermolecular forces that induces condensation below 100oC and freezing–solidification to form ice crystals below 0oC.

In the reverse process, when ice is warmed, the intermolecular forces are weakened and at 0oC the intermolecular bonds are weakened enough to allow melting to take place.

Above 0oC (evaporation), and particularly at 100oC (boiling), the intermolecular forces are weak enough for 'intact water molecules' to escape from the surface of the liquid water.

It is VERY important to realise that the chemical hydrogen–oxygen covalent bonds (O–H) in water are NOT broken and the state changes ...

solid <== freezing/melting ==> liquid <== condensing/boiling ==> gas ...

are due to the weakening of the intermolecular forces/bonds with increase in temperature OR the strengthening of the intermolecular bonds/forces decrease in temperature.

The same arguments apply to all the other small covalent molecules you will come across on your course eg methane, iodine, carbon dioxide, alkanes like hexane in petrol etc. etc.


Appendix 2. How to work out a covalent compound formula

Selected valencies of elements

The valency of an element is the combing power of its atoms.

Note that some elements can have more than one valency

Hydrogen  H (1)

Chlorine Cl and other halogens (1)

Oxygen O and sulphur S (2)

Boron B and aluminium Al (3)

Nitrogen (3, 4, 5)

Carbon C and silicon Si (4)

Phosphorus (P 3,5)

To work out a covalent compound formula by combining 'A' with 'B' the rule is

number of atom 'A' x valency of atom 'A' = number of atom 'B' x valency of atom 'B'

Three examples are shown below

'A' (valency) 'B' (valency) deduced formula of A + B
1 of carbon C (4) balances 4 of hydrogen H (1) 1 x 4 = 4 x 1 = CH4 
1 of nitrogen (3)  balances 3 of chlorine Cl (1) 1 x 3 = 3 x 1 = NCl3 
1 of carbon C (4) balances 2 of oxygen O (2) 1 x 4 = 2 x 2 = CO2 
(c) doc b The diagram on the left illustrates the three covalent examples above for

methane CH4

nitrogen trichloride NCl3

carbon dioxide CO2

 

 


top

dot and cross covalent bonding diagrams for H2 Cl2 HCl H2O NH3 CH4 O2 CO2 C2H4 N2 C2H6 CH3CH3 CH3Cl CH3OH


Revision notes information to help revise KS4 Science Additional Science Triple Award Separate Sciences GCSE/IGCSE/O level Chemistry Revision–Information Study Notes for revising for AQA GCSE Science, Edexcel GCSE Science/IGCSE Chemistry & OCR 21st Century Science, OCR Gateway Science WJEC/CBAC GCSE science–chemistry CCEA/CEA GCSE science–chemistry (and courses equal to US grades 8, 9, 10) basic aid notes for GCE Advanced Subsidiary Level AS Advanced Level A2 IB Revise AQA OCR Edexcel Salters CIE, CCEA/CEA & WJEC advanced level courses for pre–university students (equal to US grade 11 and grade 12 and Honours/honors level courses)


WHAT NEXT? and other associated Pages


top
ALL Website content copyright © Dr Phil Brown 2000–2014 All rights reserved on revision notes, images, puzzles, quizzes, worksheets, x–words etc. * Copying of website material is not permitted * chemhelp@tiscali.co.uk

The covalent bonding in simple molecules and their properties – covalently bonded elements and compounds

Teach yourself chemistry online ALPHABETICAL SITE INDEX for chemistry

Alphabetical Index for Science Pages Content A B C D E F G H I J K L M N O P Q R S T U V W X Y Z

**

SITE HELP SEARCH – ENTER SPECIFIC WORDS/FORMULA etc.