03 Covalent Bonds And Molecules

THE PARTICLES OF MATTER (PART III): COVALENT SUBSTANCES NON METALS BOND OTHER NON METALS: THE COVALENT BOND In this case, none of the atoms will give electrons to the other! Instead they share one, two or three pairs of electrons in order to achieve the complete shell. These “shared pairs” are no longer located around one of the nuclei but around both and mainly in the zone between both nuclei. A molecule is formed (see at the beginning of the chapter) Hydrogen Hydrogen atoms only need two electrons in their outer level to reach the noble gas structure of helium. Once again, the covalent bond holds the two atoms together because the pair of electrons is attracted to both nuclei. The formula of hydrogen gas, formed by hydrogen molecules is H2

Chlorine Two chlorine atoms could both achieve stable structures by sharing their single unpaired electron as in the diagram. The fact that one chlorine atom has been drawn with electrons marked as crosses and the other as dots is simply to show where all the electrons come from. There is no difference between them. The two chlorine atoms are said to be joined by a covalent bond. The reason that the two chlorine atoms stick together is that the shared pair of electrons is attracted to the nucleus of both chlorine atoms. The formula of the molecule is written Cl2. This formula shows not just the relative amount but the actual number of atoms of different elements in a molecule of the compound Hydrogen chloride Now the sharing is between two different non metals. Once again sharing takes place and a molecule is formed. The hydrogen has a helium structure, and the chlorine an argon structure. The formula for this substance is HCl. Double and triple covalent bonds are formed when two or three pairs of electrons are shared between the atoms rather than just one pair. Some Simple Molecules Containing Double Bonds Oxygen, O2 Two oxygen atoms can both achieve stable structures by sharing two pairs of electrons as in the diagram.

The double bond is shown conventionally by two lines joining the atoms. Each line represents one pair of shared electrons. Carbon dioxide, CO2

Ethene, C2H4 Ethene has a double bond between the two carbon atoms.

Simple examples for the formation of triple bonds are the nitrogen molecule in which the N atoms share three pairs of electrons, and the ethyne (acetylene) molecule with a triple C-C share. (try to sketch them both). The physical properties of molecular substances Molecules are made of fixed numbers of atoms joined together by covalent bonds, and can range from the very small (even down to single atoms, as in the noble gases) to the very large (as in polymers, proteins or even DNA). The covalent bonds holding the molecules together are very strong, but the forces among the molecules are weaker than in the other cases as it has been stated. That is why they are soluble in many solvents and show low melting points (easy to separate). Physical properties are governed by the intermolecular forces - forces attracting one molecule to its neighbours. •

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Melting and boiling points Molecular substances tend to be gases, liquids or low melting point solids, because the intermolecular forces of attraction are comparatively weak. You don't have to break any covalent bonds in order to melt or boil a molecular substance. The melting or boiling point will depend on the strength of the intermolecular forces. Solubility in water Most molecular substances are insoluble (or only very sparingly soluble) in water. Those which do dissolve often react with the water, or else are capable of forming hydrogen bonds with the water. Solubility in organic solvents Molecular substances are often soluble in organic solvents - which are molecular. Both the solute (the substance which is dissolving) and the solvent are likely to have molecules attracted to each other by weak intermolecular forces. Although these attractions will be disrupted when they mix, they are replaced by similar ones between the two different sorts of molecules. Electrical conductivity Molecular substances won't conduct electricity.

INTERMOLECULAR BONDING Intermolecular attractions are attractions between one molecule and neighbouring molecules. All molecules experience intermolecular attractions, because as we have seen at the before, covalent bonds are most of the time polar bonds. The slightly positive and the slightly negative parts of the molecules attract each other although in some cases those attractions are very weak. Even in a gas like hydrogen, H 2, if you slow the molecules down by cooling the gas, the attractions are large enough for the molecules to stick together eventually to form a liquid and then a solid. These forces are generally known as van der Waals forces INTERMOLECULAR BONDING Consider two water molecules coming close together. The δ+ hydrogen is so strongly attracted to the lone pair that it is almost as if you were beginning to form a covalent bond. It doesn't go that far, but the attraction is significantly stronger than an ordinary dipole-dipole interaction. These relatively powerful intermolecular forces are described as hydrogen bonds. The molecules which have this extra bonding are: Notice that in each of these molecules: •

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The hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant amount of positive charge. Each of the elements to which the hydrogen is attached is not only significantly negative, but also has at least one lone pair.

Hydrogen bonds have about a tenth of the strength of an average covalent bond, and are being constantly broken and reformed in liquid water. If you liken the covalent bond between the oxygen and hydrogen to a stable marriage, the hydrogen bond has "just good friends" status. On the same scale, van der Waals attractions represent mere passing acquaintances! Hydrogen bonding in Other Compounds Any molecule which has a hydrogen atom attached directly to an oxygen atom or a nitrogen is capable of hydrogen bonding. Such molecules will always have higher boiling points than similarly sized molecules which don't have an -O-H or an -N-H group. The hydrogen bonding makes the molecules "stickier", and more heat is necessary to separate them. Ethanol, CH3CH2-O-H, and methoxymethane, CH3-O-CH3, both have the same molecular formula, C2H6O. However, ethanol has a hydrogen atom attached directly to an oxygen that has two lone pairs as in a water molecule. Hydrogen bonding can occur between ethanol molecules.

The chart below shows the dramatic effect that the hydrogen bonding has on the stickiness of the ethanol molecules: ethanol (with hydrogen bonding) 78.5°C methoxymethane (without hydrogen bonding) -24.8°C

When mixed with water, molecules with OH groupings will dissolve far more easily than other covalent substances do. That is because both molecules can “interchange” hydrogen bonds. Do you guess why glucose C6H12O6 with 6 OH groups dissolves so easily in water and hexane C6H14 is almost absolutely insoluble? GIANT COVALENT STRUCTURES The so called gigantic covalent structures in which millions of millions of atoms are bonded covalently deserve a separate paragraph. We don’t have a regular molecule in this case but a giant structure. In fact the atoms form just one super-particle (if we allow the term particle to be applied in this case) and dissolving or melting it would require the breaking down of millions of chemical bonds.

The giant covalent structure of diamond Carbon has an electronic arrangement of 2,4. In diamond, each carbon shares electrons with four other carbon atoms - forming four single bonds. In the diagram some carbon atoms only seem to be forming two bonds (or even one bond), but that's not really the case. We are only showing a small bit of the whole structure. This is a giant covalent structure - it continues on and on in three dimensions. It is not a molecule, because the number of atoms joined up in a real diamond is completely variable - depending on the size of the crystal. The physical properties of diamond Diamond: •

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has a very high melting point (almost 4000°C). Very strong carbon-carbon covalent bonds have to be broken throughout the structure before melting occurs. is very hard. This is again due to the need to break very strong covalent bonds operating in 3-dimensions. doesn't conduct electricity. All the electrons are held tightly between the atoms, and aren't free to move. is insoluble in water and organic solvents. There are no possible attractions which could occur between solvent molecules and carbon atoms which could outweigh the attractions between the covalently bound carbon atoms.

The giant covalent structure of graphite Graphite has a layer structure which is quite difficult to draw convincingly in three dimensions. The diagrams below show the arrangement of the atoms in each layer, and the way the layers are spaced.

Notice that you can't really draw the side view of the layers to the same scale as the atoms in the layer without one or other part of the diagram being either very spread out or very squashed. In that case, it is important to give some idea of the distances involved. The distance between the layers is about 2.5 times the distance between the atoms within each layer. The layers, of course, extend over huge numbers of atoms - not just the few shown above. Bonding in graphite Each carbon atom uses three of its electrons to form simple bonds to its three close neighbours. That leaves a fourth electron in the bonding level. These "spare" electrons in each carbon atom become delocalised over the whole of the sheet of atoms in one layer. They are free to wander throughout the whole sheet. The atoms within a sheet are held together by strong covalent bonds - stronger, in fact, than in diamond because of the additional bonding caused by the delocalised electrons. The physical properties of graphite Graphite: •

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has a high melting point, similar to that of diamond. In order to melt graphite, it isn't enough to loosen one sheet from another. You have to break the covalent bonding throughout the whole structure. has a soft, slippery feel, and is used in pencils and as a dry lubricant for things like locks. You can think of graphite rather like a pack of cards - each card is strong, but the cards will slide over each other, or even fall off the pack altogether. When you use a pencil, sheets are rubbed off and stick to the paper. is insoluble in water and organic solvents - for the same reason that diamond is insoluble. Attractions between solvent molecules and carbon atoms will never be strong enough to overcome the strong covalent bonds in graphite. conducts electricity. The delocalised electrons are free to move throughout the sheets. If a piece of graphite is connected into a circuit, electrons can fall off one end of the sheet and be replaced with new ones at the other end.

Simple examples for the formation of triple bonds are the nitrogen molecule in which the N atoms share three pairs of electrons, and the ethyne (acetylene) molecule with a triple C-C share. The chart next page sums up what has been previously explained

NB: SUBSTANCES SHOWING BOTH COVALENT AND IONIC BONDS BEHAVE AS IONIC SUBSTANCES Class

Formed by

Structure (scheme)

Covalent Molecules

Melting point

Solubility

Conducts electricity

<300º C

Soluble en solvents And not in Never water (generally)

>1500º C

Insoluble in Only any liquid graphite conducts

Example

Oil, petrol

Macromolecules Covalent (macro)

Graphite, quartz

THE SHAPE OF MOLECULES The electron pair repulsion theory The shape of a molecule or ion is governed by the arrangement of the outer shell’s electron pairs around the central atom. All you need to do is to work out how many electron pairs there are at the bonding level, and then arrange them to produce the minimum amount of repulsion between them. You have to include both bonding pairs and lone pairs. How to work out the number of electron pairs You can do this by drawing dots-and-crosses pictures.  First you need to work out how many groups there are bonded around the central atom.  Now work out how many lone (non bonding) pairs of electrons there are. These pairs will count as a group. If there are multiple bonds, the second or third pairs of electrons forming the bond will not count  Finally, work out the shape. Arrange the groups and lone electron pairs in space to minimise repulsions. How this is done will become clear in the examples which follow.  Lone pairs being more “free to move” occupy more space and will push other groups backwards so that the angle between them and other groups is slightly wider. Four electron pairs around the central atom There are lots of examples of this. The simplest is methane, CH4. Carbon is in group 4, and so has 4 outer electrons. It is forming 4 bonds to hydrogen atoms, adding another 4 electrons - 8 altogether, in 4 pairs. Because it is forming 4 bonds, these must all be bonding pairs. Four electron pairs arrange themselves in space in what is called a tetrahedral arrangement. A

tetrahedron is a regular triangularly-based pyramid. The carbon atom would be at the centre and the hydrogen atoms at the four corners. All the bond angles are 109.5°. Other examples with four electron pairs around the central atom Ammonia, NH3 Nitrogen is in group 5 and so has 5 outer electrons. Each of the 3 hydrogen atoms is adding another electron to the nitrogen's outer level, making a total of 8 electrons in 4 pairs. The electron pairs arrange themselves in a tetrahedral fashion as in methane. But because the nitrogen is only forming 3 bonds, one of the pairs must be a lone pair. lone pairs occupy more space and will push other groups backwards so that the angle between them and other groups is slightly wider. 

Remember this: Although the electron pair arrangement is tetrahedral, when you describe the shape, you only take notice of the atoms. Ammonia is pyramidal - like a pyramid with the three hydrogen atoms at the base and the nitrogen at the top. Water, H2O Following the same logic as before, you will find that the oxygen has four pairs of electrons, two of which are lone pairs. These will again take up a tetrahedral arrangement. This time the bond angle closes slightly more to 104°, because of the repulsion of the two lone pairs. The shape isn't described as tetrahedral, because we only "see" the oxygen and the hydrogens - not the lone pairs. Water is described as bent or V-shaped. THE SHAPE OF MOLECULES WITH THREE GROUPS AROUND THE CENTRAL ATOM The simple cases of this would be BF3 or BCl3. Boron is in group 3, so starts off with 3 electrons. It is forming 3 bonds, adding another 3 electrons. There is no charge, so the total is 6 electrons - in 3 pairs. Because it is forming 3 bonds there can be no lone pairs. The 3 pairs arrange themselves as far apart as possible. They all lie in one plane at 120° to each other. The arrangement is called trigonal planar. In the diagram, the other electrons on the fluorines have been left out because they are irrelevant. THE SHAPE OF MOLECULES WITH TWO GROUPS AROUND THE CENTRAL ATOM Two groups around the central atom As we have already seen in carbon dioxide C atom forms two double bonds. As the bonding pairs to the oxygen atoms will try to be as far as possible

from each other, they arrange themselves at 180° to each other. The molecule is described as being linear. HCN (hydrogen cyanide) belongs to this group

PROBLEMS 1- Carbon dioxide is a compound in which carbon forms double bonds to oxygen (shares two pairs of electrons). Write the cross and dot diagram for carbon dioxide. (Cross and dot : see the diagrams of Cl2, HCl and H2) 2- Write a cross-and-dot diagram for the ammonia molecule NH3 3- Silicon is an element that melts at 1.410 °C and a poor conductor. a- What can you say about its structure? Explain b- What can you predict about its solubility both in water and in organic solvents? Explain. 4- The chart shows some properties of tour substances A, B, C and D. Substance Melting point

Electric Solubility in conductivity water Solid Melted A 961 °C Good Good insoluble B 1610 °C Poor Poor insoluble C 776 °C Poor Good very soluble D 37 °C Poor Poor insoluble Using the information in the text (see chart) decide which of them is: a- Eicosane? b- Silver? c- Quartz? d- Potassium chloride? e- Should eicosane’s formula be C20H42 or rather SrO?

Solubility solvents insoluble insoluble insoluble soluble

5- Explain the following facts a- Graphite conducts electricity but diamond does not. b- You can write with a graphite pencil. c- Glucose C6H12O6 melts at 146 ºC and NaCl at 801 d- H2O is a liquid (boils at 100 ºC) and H2S is a gas (boils at -61 ºC) 6- The following table shows the melting point (in K) of the elements of the 3rd period. Explain the trend.

7- Which factors affect the solubility of molecular substances in water? 8- Which of each pair can be predicted to have a higher solubility in water?

in

a- CH3-CH2-O-CH2-CH3

or

CH3-CH2-CH2-CH2-OH

b- CH3-CH2-CH2-CH2-CH3

or

CH3-CH2-O-CH2-CH3