BONDING - Barnard Castle School€¦ · Web viewHydrogen bonding explains the unusually high melting and boiling points of molecules such as water. With an MR of only 18, water

BONDING

Name: ………………………………………….

A chemical bond is a method of holding atoms or ions together. You should already be familiar with the three types of chemical bond: ionic, covalent and metallic.

1. Ionic Bonding Metals, since they are on the left hand side of the Periodic Table, tend

to lose electrons to gain a stable full outer shell. The tendency for an atom to lose electrons is greatest at the bottom

left of the Periodic Table. Non-metals, at the opposite end of the Periodic Table, tend to gain

electrons. The tendency for a non-metal to gain electrons is greatest at the top

right of the Periodic Table.

If a metal and a non-metal (e.g. sodium and chlorine) come into contact, there is a tendency for electrons to be transferred from the metal to the non-metal.

The metal atom loses electrons and becomes a positively charged ion (cation). Draw Na+

The non-metal gains electrons and becomes a negatively charged ion (an anion). Draw Cl-

These opposite charges attract and the ions are held together by strong electrostatic forces of attraction. The formation of an ionic bond is represented by a dot and cross diagram:

Exercise 1: Draw diagrams to show the ionic bonding in

(a) lithium fluoride

TOPIC 12.3: BONDING 1

(b) magnesium chloride

(c) calcium oxide

(d) sodium oxide


Ionic Formulae

You should be able to work out the charge on some ions from its position on periodic table. You should also know the formulae of some very common compound ions (see table below). It is then possible to work out the correct formula for an ionic compound.

Ion Formulahydroxide OH-

sulphate SO42-

nitrate NO3-

carbonate CO32-

ammonium NH4+

Eg. What is the formula of magnesium nitrate?

Magnesium is in group 2 of the periodic table and will therefore forms Mg2+

ions. The nitrate ion is NO3- and when these two ions combine to form a neutral compound, they do so in a ratio of 1:2. The formula is therefore Mg(NO3)2. You should remember from GCSE that if you need to show the presence of more than 1 compound ion, we use brackets.

What is the formula of iron (II) hydroxide?

Iron is a transition metal and forms ions with different charges, however, the roman numerals tell us which ion we are dealing with. We will see later that this number is called the oxidation number of the iron.

If we have Fe2+ and our negative ion is OH-, these ions will combine in a ratio of 2:1 to form a neutral compound and therefore the formula of iron (II) hydroxide is Fe(OH)2.

Try a few more

compound formula compound formula

1 silver iodide ammonium chloride

2 sodium carbonate calcium bromide

3 potassium sulphate magnesium nitrate

4 ammonium sulphate calcium hydroxide

5 aluminium nitrate lead hydroxide

6 iron(II) nitrate ammonium carbonate

7 zinc iodide magnesium iodide

8 copper(II) hydroxide aluminium oxide

9 iron(III) oxide iron(III) sulphate

The Structure of Ionic Solids

A description of the bonding in a compound describes the way two atoms or ions are held together, as above.


The structure of a solid describes how the many particles that make up that solid are arranged and held together.

A solid with a regular shape which contains particles arranged in a regular structure is called a crystal.

The physical properties of that solid crystal are determined by the bonds and forces that hold the crystal together.

In an ionic solid, the individual particles are positive and negative ions (cations and anions), arranged in a regular array (a lattice) e.g. sodium chloride, NaCl:

The ions in a sodium chloride crystal are arranged in a 3-dimensional giant ionic lattice. Each sodium ion has a coordination number of 6 (i.e. each Na+ is surrounded by 6 Cl-) and each chloride ion has a coordination number of 6. The sodium chloride giant ionic lattice can thus be described as a 6:6 lattice or a face centred cubic lattice (because the resulting crystals are cubes).

Each sodium ion is in contact with six chloride ions; each chloride ion is in contact with six sodium ions.


Na+ ion

Cl- ion

In reality, the ions are touching, as shown in the following diagrams.

Other ionic compounds could have different lattice structures. The structure depends on the relative size of the anions and cations. For example, caesium chloride has a body centred cubic lattice and is described as an 8:8 lattice.


single layer

Explaining the physical properties of ionic solids

You should be familiar with the properties of ionic solids. These properties can be explained by reference to the structure of the compounds.

Exercise 2: Explain the following properties of ionic solids in terms of their structure.

1. Ionic compounds have high melting and boiling points

2. Ionic compounds are hard

3. Ionic compounds are brittle

4. Ionic compounds are non-conductors of electricity when solid, but will conduct when molten (in the liquid state) or in solution


Polarisation of Ions(Ionic Compounds with Covalent Character)

An ionic compound forms when electron transfer takes place between a metal and a non-metal to produce a cation and an anion, which are then held together by strong electrostatic forces.

If the electron is completely transferred, and the ions formed are perfectly spherical, the bonds are perfectly ionic. This is more likely to happen between metals at the bottom left of the periodic Table, and non-metals at the top right e.g. caesium fluoride

Cs+ F-

However, if the positive ion (cation) formed is very small and/or highly charged, and the negative ion (anion) is large and/or not highly charged the electron cloud around the anion is distorted. It is no longer spherical. Because the electron density of the anion is now partially localised between the two nuclei, the resulting ionic compound is said to have covalent character e.g. lithium iodide:

Li+ I-

The type of bonding in lithium iodide is said to be ‘ionic with covalent character’.

A cation which is small and/or highly charged is said to be polarising. It polarises the anion, as shown by Li+ above. The greater the charge density (i.e. a high charge on a small cation) the greater the polarising power of the cation.

An anion which is large is said to be polarisable i.e. it can be polarised, as in the case of I- above. The larger anions are more easily polarised by a cation.


Which ionic compounds are likely to display covalent character?

The Metal CationCation size increases down a Group, so the most polarising cations are likely to be at the top of a Group.

Cation size decreases across a Period (from Group I to Group III), so the polarising power increases from left to right on the Periodic Table.

Thus, metal cations from the top of Groups, and furthest across the Period on the Periodic Table are likely to be the most polarising e.g. Be2+ and Al3+, and are more likely to form ionic compounds with covalent character.

The Non-Metal AnionAnion size increases down a Group, so the most polarisable (the most easily distorted) anions are likely to be near the bottom of a Group.

Anion size decreases across a Period (from Group V to Group VII), so the most polarisable anions are P3-, then S2- and then Cl-.

Thus, non-metal anions from the bottom of Groups, and furthest across the Period on the Periodic Table are likely to be the most polarisable (Their spherical shape is most easily distorted) e.g. I-, and are more likely to form ionic compounds with covalent character.

The effect of covalent character on physical properties

An ionic compound with covalent character is likely to show properties which resemble those typical of covalent compounds.

For example, an ionic compound which displays covalent character is likely to have a lower melting point than expected. It is likely to be softer, less brittle and show reduced electrical conductivity when molten or in solution.


Exercise 3: Circle the ionic compound in each of the pairs below that exhibits the most covalent character. Give a reason for your choice.

Ionic Compounds Reason

LiCl and NaCl

LiCl and LiI

NaCl and MgCl2

CaO and CaS

MgCl2 and AlCl3

KBr and KF


Special cases: Aluminium Chloride and Beryllium Chloride

The Al3+ ion has such a small size and high charge that it is highly polarising. It has a very high charge density.

When aluminium forms a compound with oxygen, Al2O3, the ionic bonding has only very slight covalent character because, despite the high charge density of the Al3+, the O2- ion is not very polarisable.

When aluminium bonds with chlorine, however, the resulting chloride ions are polarised to such an extent, and the bond has such a large amount of covalent character, that any ionic description of the bonding is useless.

The bonding in AlCl3 is COVALENT, and AlCl3 is a molecule, despite the compound being formed from a metal and a non-metal. It can be drawn as follows:

AlCl3 molecule:

The same argument and reasoning would apply to BeCl2, which is also a covalent compound made of a metal and a non-metal.

Two AlCl3 molecules form an Al2Cl6 dimer – see notes on covalent bonding.

Across Period 3 from Na to Si, the type of bonding in the chlorides changes as follows:

NaCl MgCl2 AlCl3 SiCl4ionic ionic with a little

covalent character

covalent covalent


Melting an Ionic Solid

In an ionic crystal, e.g. sodium chloride, the ions, formed by electron transfer, are arranged in a regular array and are all held together by strong electrostatic forces. The ions vibrate around a fixed point. This description of a solid should be familiar from your GCSE course.

If the system is heated (i.e. it takes in energy), the ions vibrate faster and faster until the melting point is reached.

At the melting point, the particles have taken in enough energy to partially overcome the strong electrostatic forces holding them in a fixed position. They can now move around each other. The electrostatic forces still act, however, to keep the particles close together in the liquid (molten) state, but the ions are free to move around. Again, the motion of the ions in the molten state should be familiar from your GCSE description of the motion of particles in a liquid.

Because the melting process requires (takes in) energy, it is described as ENDOTHERMIC.

The large amount of energy required to partially overcome the many strong electrostatic forces is the reason that ionic solids have high melting points.

If the liquid (molten) ionic compound continues to be heated, the ions take in more energy and move faster and faster until the boiling point is reached. At this point, the ions possess enough energy to almost completely overcome all electrostatic forces, and the ions move away from each other in random, chaotic motion. It is now a gas.


2. Covalent Bonding Non-metal atoms need to gain electrons to achieve a full outer shell. When two non-metals bond, neither one will give up an electron, so the

atoms share a pair of electrons. This means that each atom has a full outer shell, despite some of the

electrons in the outer shell being shared with another atom. A covalent bond is defined as a shared pair of electrons.

A covalent bond can be represented by a dot and cross diagram. Only the outer, bonding electrons need to be shown.

Hydrogen fluoride

The dot and cross diagram can be simplified, and the covalent bond shown as a single line between two atoms. Every time you see a diagram with a line of this sort, it means the atoms are covalently bonded together.

Hydrogen fluoride Chlorine


Exercise 4: Draw similar, simplified diagrams for the following molecules. You may find it easier to first sketch the dot and cross diagrams. On your finished diagrams, indicate the number of bonding and non-bonding pairs (known as lone pairs) of electrons.

(a) methane (CH4)

(b) ammonia (NH3)

(c) water (H2O)

(d) hydrogen sulphide (H2S)


Dative Covalent (Coordinate) Bonding

The covalent bonds studied so far have been formed between two atoms, each of which donates a single electron to make a shared electron pair.

However, both electrons may originate from the same atom, and be shared with another atom.

A dative covalent bond (also known as a coordinate bond) is defined as a shared pair of electrons where both electrons have originated from one atom.

For this to happen, one species must have a lone pair of electrons which it can donate to another atom in order for the electrons to be shared. This atom is the electron pair donor. The other atom acts as an electron pair acceptor.

A dative covalent bond is identical to a normal covalent bond except in the way it forms.

The formation of a dative covalent bond can be represented as shown below. The lines represent a shared pair of electrons, and the two crosses, xx, represent the lone pairs.

Boron trifluoride gas and ammonia gas react together readily to give a white solid.

The boron atom acts as an electron pair acceptor (so BF3 is known as a Lewis acid – an electron pair acceptor) because it is electron deficient.

The nitrogen atom in ammonia acts as an electron pair donor (so NH3 is known as a Lewis base – an electron pair donor).


Exercise 5: When ammonia and hydrochloric acid react they produce a salt called ammonium chloride. Draw a diagram to show how ammonia can react with the H+ ion in the acid.

Exercise 6: Two AlCl3 molecules form a dimer of Al2Cl6. Draw a diagram to show how this dimer forms.

Exercise 7: When copper(II)sulphate solution is a blue colour due to the presence of the [Cu(H2O)6]2+ complex ion. This forms between Cu2+ ions and water molecules. Draw a diagram to show how this complex ion forms.


Electronegativity and Bond Polarity

In a covalent bond, a pair of electrons is shared between two atoms. Some atoms attract the bonding pair of electrons in a covalent bond

more than others. The ability of an atom to attract the bonding pair of electrons in a

covalent bond is known as the electronegativity of the atom.

What factors do you think might affect how electronegative an atom is?

Linus Pauling assigned electronegativity values for each element. The scale was arbitrary, and ranged from 0 (not electronegative, i.e. does not attract the electron pair in a covalent bond) to 4 (the most electronegative – attracts the electron pair in a covalent bond very strongly). This is shown for selected elements below.

2.1H

1.0Li

1.5Be

2.0B

2.5C

3.0N

3.5O

4.0F

0.9Na

1.2Mg

1.5Al

1.8Si

2.1P

2.5S

3.0Cl

0.8K

1.0Ca

2.8Br

0.8Rb

1.0Sr

2.5I

0.7Cs

0.9Ba

How does electronegativity vary across a Period? Why?

How does electronegativity vary down a Group? Why?

Which is the most electronegative atom? Why?

Why are the Noble Gases not included in this table?If a covalent bond forms between atoms with DIFFERENT electronegativities, the electrons are not shared equally. The shared pair is held closer to the more electronegative atom. This type of covalent bond is known as a polar covalent bond, and it results from the differing electronegativities of the two atoms involved.

The atom which is the most electronegative of the two will have the greater share of the bonding pair of electrons. Because electrons are negatively charged, the more electronegative atom gains a very slight negative charge TOPIC 12.3: BONDING 16

and is said to be electron rich. This is shown using a small - symbol next to the atom.

Similarly, the atom which has the lowest share of the electron pair is said to be electron deficient and has a slight positive charge. This is shown using a small + symbol next to the atom.

Bond polarisation is caused by the unequal sharing of electrons due to the differing electronegativities of the atoms involved. As a result of bond polarisation, the covalent compound may have some degree of ionic character, because the + and - charges represent slight charges.

Exercise 8: Use the electronegativity values to deduce which atoms are electron deficient and which are electron rich. Indicate, using + and - symbols on the appropriate atoms, the bond polarisation of any one of the bonds in the molecules shown below. If the bond is not polar, write ‘non-polar’ underneath the structure.

(a)

(b)

(c)

(d)

(d)


Calculating Percentage Ionic Character

Linus Pauling estimated the amount of partial ionic character in a covalent bond in a binary compound (a compound consisting of 2 atoms). He calculated the electronegativity difference of the two atoms involved in the bond, and thus estimated the percentage ionic character of the resulting molecule.

Electronegativity Difference

% Ionic Character

Electronegativity

Difference

% Ionic Character

0.1 0.5 1.7 510.2 1 1.8 550.3 2 1.9 590.4 4 2.0 630.5 6 2.1 670.6 9 2.2 700.7 12 2.3 740.8 15 2.4 760.9 19 2.5 791.0 22 2.6 821.1 26 2.7 841.2 30 2.8 861.3 34 2.9 881.4 39 3.0 891.5 43 3.1 911.6 47 3.2 92

Exercise 9: Use the electronegativity values to deduce the percentage ionic character in the following binary compounds:

Compound % Ionic CharacterChlorine, Cl2

Hydrogen chloride, HCl

Sodium chloride, NaCl

Magnesium oxide, MgO

Caesium fluoride, CsF


Symmetry and bond polarity

In symmetrical molecules with polar covalent bonds, the bond polarities can ‘cancel out’. The results is a molecule which, overall, is non-polar but which contains polar bonds.

- + -

O=C=O

Despite the two polar C=O bonds, CO2 is a non-polar molecule.

Exercise 10: Indicate if the following molecules are overall polar or non-polar.

(a) HCl (b) NH3

(c) H2C=O (d) H2O

(e) BeH2 (f) AlF3


Water is a polar molecule:

Exercise 11: The structure of chloromethane and tetrachloromethane are shown below. For each carbon-chlorine bond, show the bond polarity by using the symbols - and + to indicate the electron rich and electron deficient atoms. Draw a circle around the molecule that does not possess an overall dipole.

Cl Cl

H Cl H H Cl Cl

chloromethane tetrachloromethane

Both molecules have 4 bonds arranged in a tetrahedral fashion around the carbon atom.

Exercise 12: Carbon dioxide (CO2) contains polar C=O bonds but it is a symmetrical molecule so it shows no overall dipole. Sulphur dioxide (SO2) does have an overall dipole. Draw the structure of an SO2 molecule, which could be consistent with the information given. Hint: carbon has 4 outer electrons and sulphur has 6.


The Structure of Covalent Solids

There are two main types of covalent structures: Giant Covalent and Simple Molecular. The two types have particles arranged and held very differently in the solid state, and so have very different physical properties (e.g. melting point, conduction of electricity, hardness, malleability).

(1) Giant Covalent Structures

Giant covalent structures have of a regular array (lattice) of atoms held by many strong covalent bonds. Examples are diamond and graphite (see below).

The giant covalent structures are very strong. As a consequence the melting points are very high because a large amount of energy is required to break the many strong covalent bonds.

They are insoluble in water (a polar solvent) and organic solvents (non-polar solvents) because the solvent cannot break down the lattice structure.

Giant covalent structures do not conduct electricity because there are no free/mobile electrons (graphite is an exception and can conduct electricity – see below).

Examples of giant covalent structures are diamond, graphite and silicon dioxide:

Diamond structure

Silicondioxide structure

Graphite structure


Examples of Giant Covalent Structures:

Diamond, silicon and silicon dioxide (SiO2)

Graphite

3-dimensional tetrahedral lattice hexagons of carbon atoms arranged in layers (sheets)

Each carbon atom has a coordination number of 4

Each carbon has a coordination number of 3

All carbons joined to others by strong covalent bonds

Each carbon is covalently bonded to 3 others.

Diamond cannot conduct electricity because there are no delocalised (mobile) electrons

The remaining electron from the carbon is delocalised between the layers (sheets). These delocalised electrons mean that graphite can conduct electricity.

All the atoms are held by 4 very strong covalent bonds. This makes diamond very hard.

The layers (sheets) are held together by weak Van der Waals’ forces (see later). This means that the structure is easily broken along the layers (the layers can slide over each other).


Exercise 11: Explain the following observations

(a) diamond and graphite have high melting points.

(b) graphite will conduct electricity but diamond will not.

(c) graphite is soft, but diamond and silicon are very hard.


(2) Simple Molecular Structures

Simple molecular structures are made up of molecules held together by weak intermolecular forces.

A molecule is defined as one or more atoms held together by covalent bonds. Examples are Cl2, H2O, HCl, P4 and S8 (shown below).

A P4 molecule A S8 molecule

Intermolecular forces are forces which act between molecules to hold them together. They are weak forces. There are several different types of intermolecular forces that can act between molecules.

In order to understand the structure and properties of simple molecular crystals, it is essential that the intermolecular forces that hold the molecules in the solid are fully understood.

There are three types of intermolecular forces you need to consider:

1. van der Waals forces2. permanent dipole-dipole interactions3. hydrogen bonding (which is not a bond but an intermolecular force

for A level purposes!!)

All intermolecular forces are weak forces – nowhere near as strong as covalent, ionic or metallic bonds.

The three types, however, have different strengths. They are listed above in order of increasing strength, so hydrogen bonding is the strongest intermolecular force, then permanent dipole-dipole interactions. The weakest intermolecular forces are van der Waals forces.

All species, even the atoms of the Noble Gases, are attracted to each other by van der Waals forces.

Only polar molecules contain permanent dipole-dipole forces in addition to van der Waals.

The only molecules that contain hydrogen bonding are those that have a hydrogen atom covalently bonded to a N, O or F atom within the molecule.

Intermolecular forces

Van der Waals forces


Forces of attraction exists between all molecules. These weakest of these forces are known as van der Waals forces.

They originate when two molecules approach each other. The electron cloud within the molecule can become repelled at one end by repulsion from the electron cloud of the second molecule. This causes a temporary dipole, which in turn causes (induces) an opposite dipole on the adjacent molecule. The leads to very weak electrostatic attractions between molecules.

The greater the size of the molecule, the greater the number of electrons and so the greater the probability that electrons will be temporarily localised at one end of the molecule. Larger molecules also have a greater surface area for the van der Waals forces to act over.

Thus, the larger the molecule, the stronger the van der Waals forces between the molecules. Stronger forces between molecules mean that the melting and boiling point of the solid is higher.

The more branched an organic molecule is, the lower the surface contact area and thus the weaker the van der Waals forces and the lower the boiling point.

Van der Waals forces are the only intermolecular forces that act between non-polar molecules, and are the forces that hold molecules together in a solid and a liquid. There are even very weak van der Waals forces between gas atoms or molecules.


Permanent dipole-dipole attractions

We have already studied some molecules that are polar. Bond polarities arise when the electronegativities of the two atoms that form the covalent bond are different. Polar molecules are those which are non-symmetrical and which contain polar covalent bonds. Examples of polar molecule are HCl and propanone (shown below).

When two polar molecules arrange themselves in a solid, opposite charges align to give small, very weak electrostatic attractions between the molecules. This is known as a permanent dipole-dipole attraction, and it helps to hold the simple molecular structures of polar molecules together.

Van der Waals forces would also act between polar molecules, but they are weaker than permanent dipole-dipole attractions. This means that simple molecular solids containing molecules held together by both van der Waals and permanent dipole-dipole attractions are likely to have higher melting and boiling points than those with a similar mass but only van der Waals forces acting between molecules.

e.g. butane, C4H10, with an MR of 58, has a boiling point of 0oC

propanone, C3H6O, also with an MR of 58, has a boiling point of 56oC


Hydrogen bonding (it’s not a bond, it’s an intermolecular force!)

A hydrogen bond can be thought of as a very extreme example of a permanent dipole-dipole attraction.

A hydrogen bond is the force of attraction between a very electron deficient hydrogen atom and the lone pair from an atom on a neighbouring molecule.

In order that the hydrogen atom is sufficiently electron deficient, it must be directly covalently bonded to a small electronegative atom. The only three atoms which are both small enough and electronegative enough to cause this are NITROGEN, OXYGEN and FLUORINE. The diagram below shows the hydrogen bonding between two water molecules.

x + -

x

The reason hydrogen is a special case is that it has no inner shells of electrons. When hydrogen is directly bonded to N, O or F, the electron density is drawn away from the hydrogen, leaving the nucleus exposed. The nucleus of the poorly shielded hydrogen atom attracts the lone pair of electrons from an atom on the next molecule. There is some resemblance to a covalent bond – although a hydrogen bond is nowhere near as strong as a covalent bond (approximately 5-10% the strength). This explains why the hydrogen bond is stronger than a permanent dipole-dipole attraction.

Hydrogen bonding explains the unusually high melting and boiling points of molecules such as water. With an MR of only 18, water should really be a gas at room temperature. If it wasn’t for the hydrogen bonding increasing the melting and boiling point so that it is a liquid at room temperature, life on earth would not exist as we know it today!


Exercise 12: Explain why the boiling points of the halogens increase from fluorine (F2) to iodine (I2).

Exercise 13: Predict, with a brief reason, which organic compound in each pair would have the higher boiling point. Circle the compound with the highest boiling points (NB All hydrocarbons are considered to be non-polar molecules).

Organic Compound Reason

and

and

and

and

and


Exercise 14: State the name of the strongest intermolecular force present between molecules of each of the species below. (Lone pairs have not been included to simplify the diagrams).


Exercise 15: Sketch 2 graphs on the same axes to show the boiling points of the hydrides of (a) Group VI and (b) Group VII using the data below:

Group VI Hydride

Boiling Point/K

Group VII Hydride

Boiling Point /K

H2O 373 HF 293H2S 212 HCl 188H2Se 232 HBr 206H2Te 271 HI 238

Boiling Point /K

Hydride

Explain why there is a regular increase in the boiling points of the hydrides from the 2nd to the 4th member of the group.

Explain why the first hydride in each group shows an exceptionally high boiling point when compared to the rest of the compounds.

Draw a similar graph below and draw the predicted shapes of the graphs for the group V hydrides (NH3 to SbH3) and group IV hydrides (CH4 and SnH4).


Exercise 16: Draw diagrams to show the formation of one hydrogen bond between (a) molecules of hydrogen fluoride and (b) molecules of ammonia. Show all the electron deficient and electron rich atoms (using + and - symbols) and include all lone pairs.

(a) Hydrogen fluoride, HF

(b) Ammonia, NH3


The Structure of Ice

Now that we have an understanding of the forces that can hold molecules together in simple molecular structures, we can look at some examples.

Water has extensive hydrogen bonding between molecules and it can form a three dimensional lattice. It is possible for 4 hydrogen bonds to form between each water molecule – two from each lone pair, and two to each hydrogen atom per molecule.

The diagram below shows how it is possible for each water molecule to connect to 4 other molecules through hydrogen bonds (note the loose similarity to a diamond type structure – although in the case of water the structure is held together by hydrogen bonding which is much weaker than the covalent bonds that hold the giant covalent diamond structure together).

The structure above explains why ice has a lower density than water (i.e. why it expands on freezing, and why ice floats on water) at 0oC.

The hydrogen bonds hold ice in an open structure. When ice melts, the lattice breaks up, the water molecules pack closer together and the maximum density occurs at 4oC.

As the temperature increases, the hydrogen bonds are overcome, more spaces develop between the molecules so the density decreases.


The Structure of Iodine

In iodine, the I2 molecules (consisting of two iodine atoms held together by a covalent bond) are held together in a three dimensional lattice by weak van der Waals forces. They have a ‘herringbone’ pattern.

The solid is only held be very weak intermolecular forces, and so on gentle heating it will sublime.

Molecular lattices do not conduct electricity as there are no charged particles (either electrons or ions) free to move through the lattice.

Many molecular solids (e.g. iodine, phosphorus and sulphur) are insoluble in water. Can you think why this would be? What type of solvents would dissolve these types of solids?

What type of substances would dissolve in water?


Shapes of Molecules

Molecules are different shapes because electron pairs will maximise their distance apart in order to minimise repulsion between them.

Basic Shapes of Molecules

These basic shapes should all be learned. You should be able to draw accurate diagrams, name the shapes and label the bond angles.

No. of bonding electrons

Name of shape

Bond Angles

Example

Diagram

2 linear 180o BeCl2

3 trigonal planar

120o BCl3

4 tetrahedral 109.5o CH4

5 trigonal bipyramidal

120o and 90o PF5

6 octahedral 90o SF6


Exercise 17: Practice your diagrams below for CH4, PF5 and SF6. State the name of the structure underneath, and the bond angles.

Origin of names of shapes of molecules

Molecules are named after the three dimensional shapes they would form if their outer atoms were positioned at the vertices e.g.

Methane, CH4, is tetrahedral:

PF5 is trigonal bipyramidal:

SF6 is octahedral:


The presence of lone pairs

If not all the pairs of electrons around the central atom are bonding electrons, the shape can be affected.

Each electron pair occupies a volume of space, which can be regarded as a cloud of electron density. A lone pair of electrons repels more than a bonding pair of electrons because it takes up more space at the surface of the atoms, and is under the influence of only one nucleus.

The order of repulsion is:

lone pair-lone pair > lone pair-bond pair > bond pair-bond pair

Ammonia, NH3, and water, H2O, have 4 electron pairs around the central atom, but some are lone pairs. Ammonia has 3 bonding pairs and 1 lone pair around the nitrogen. Water has 2 bonding pairs and 2 lone pairs around the oxygen.

The shape of both molecules is thus based on tetrahedral (because they have 4 electron pairs around the central atom), but lone pairs occupy some of the positions around the atom. The name of the shape is determined by the position of the atoms and not the lone pairs, thus the name changes.

The bond angle is also affected, since lone pairs repel more than bonding pairs.

Molecule Diagram Shape Bond Angle

Methane, CH4

tetrahedral 109.5o


. .

Ammonia, NH3

pyramidal 107o

Water, H2O bent planar 104.5o


. .

.

.. .

Determining the Shape of Molecules

The arrangement in space of single covalent bonds around the central atom (and thus the shape of the molecule) can be predicted by considering the number of electrons in the outer shell of the central atom, and the number of bonds that it forms.

For a given molecule, the shape can be predicted using the rules below:

1. Count the total number of electrons in the outer shell of the central atom.

2. If the species is an ion, add 1 for each negative charge and subtract 1 for each positive charge.

3. Add 1 for each atom bonded to the central atom.

4. Total the number of electrons.

5. Divide by 2 to determine the total number of electron pairs. This tells you the shape your molecule is based on.

6. Count the number of bonded atoms. This tells you the number of bonding pairs of electrons. Any remaining bonds are lone pairs, and their effects must be considered.

7. State the name of the shape, draw it and show any bond angles.

Example: Draw and name the shape of PCl3, indicating any bond angles on your diagram. Following the steps as outlined above:

1. 5 electrons around P (Group 5)2. 0 to add or subtract – it’s not an ion3. 3 bonding Cl atoms4. Total number of electrons = 5 + 0 + 3 = 8 electrons5. 8 ÷ 2 = 4 PAIRS OF ELECTRONS

(so at this stage you know that the shape is based on tetrahedral – like methane)

6. 3 bonded atoms so 3 bonding pairs. The remaining pair is a lone pair.

(so I can draw a basic tetrahedral shape, put three Cl atoms at three of the positions, and show a lone pair in the remaining position.)

7. Sketch the shape, state the name of the shape and draw on the bond angles (remembering to adjust the shape name and bond angle due to the presence of the lone pair):


Exercise 18: Predicting the shapes of molecules and ions. Draw and name the shapes of the following molecules and ions. Include all the lone pairs, where appropriate, and label the bond angles. Do your working on a separate piece of paper.

BeCl2 NH3 CCl4 NH2-

H3O+ AlF63- XeF4 SiH4

AlH4- H2S PF5 NH4+

I3- (hint: use one I as the central atom)

SF4 BrF3 AlI3


3. Metallic BondingMetallic bonding, as the name suggests, occurs between metal atoms. Metals tend to want to lose electrons, and they have the ability to lose their outermost electrons to form cations.

A metal can be thought of as a giant lattice of close-packed metal cations surrounded by a sea of delocalised electrons. These delocalised electrons come from the outer shell of the metal ion, and are free to move through the lattice.

A metallic bond is the electrostatic force of attraction that two neighbouring cations have for the delocalised electrons between them.

Diagram to show metallic bonding

Close packed metal ions in the Giant Metallic Solid:

Metals have high melting points due to the strong electrostatic attraction between the metal cations and the sea of delocalised electrons.

Conduction of electricity is possible due to the delocalised (and therefore mobile) electrons, which can move freely through the structure.

The metallic bond strength varies between metals. This can be explained by the charge on the cation, the size of the cation and the number of delocalised electrons holding the cations together. The highest melting point will occur when the cation size is small, the cation charge is large and the number of delocalised electrons is large.

Exercise 19: State then explain the trend in melting points(a) across Period 3 (sodium to aluminium)(b) down Group 1 (sodium to rubidium)


+ + ++

++

++

+ ++

++

+ +free electron

metal ion

STATES OF MATTER

SOLIDSIn a solid, particles (which may be atoms, ions or molecules) are arranged in a definite, ordered regular pattern in three dimensions, and the particles are very tightly packed. This arrangement manifests itself on a macroscopic scale in the crystalline structure of solids. At absolute zero, the particles are stationary, but as the solid warms up, the particles vibrate about a mean position.

As the temperature increases, the vibrations increase in amplitude and eventually become so great that the attractive forces between the particles in the solid are no longer sufficient to hold the structure together. At this point melting occurs.

A solid has a definite shape and a definite volume.

The different types of solid structure: ionic, simple molecular, macromolecular and metallic, together with their associated properties have been discussed in detail in an earlier section.

LIQUIDSThe liquid state is the least understood of the three states of matter.

In a liquid, the tendency of particles to stick together because of their mutual attraction outweighs the tendency to remain apart because of their thermal energy. The thermal energy is however too great to allow the particles to occupy fixed positions.Thus, the particles in a liquid are still fairly tightly packed together but lack the highly ordered arrangement of a solid. Liquids are often described as having short range order and long range disorder. This means that over a small region (1 to

2 nm) there is order comparable to that in a crystal. On a larger scale there is much disorder caused by the presence of ‘holes’ in the structure. These holes allow individual particles to have translational movement, which means that particles can move past each other. Structures similar to those in a solid, but only partially complete, exist temporarily, constantly breaking and reforming in a different way as particles move through the liquid.

The structure of a liquid manifests itself on a macroscopic scale by it having a definite volume but no definite shape.

The particles in a liquid have a range of kinetic energies. Some of the faster moving particles are able to overcome the attractive forces in the liquid and escape into the space above the liquid, forming a vapour. This vapour exerts a pressure known as the vapour pressure.


That a liquid more closely resembles the solid state than the gaseous state is indicated by the following observations:

DensityThere is only a small difference between the density of a solid and that of the liquid formed from it.

Energy DifferencesMuch more energy is needed to convert a liquid to a gas than a solid to a liquid.e.g. H2O(l) H2O(g) H = +40.6 kJ.mol-1 H2O(s) H2O(l) H = +6.0 kJ.mol-1

GASESGases are devoid of structure. Since the energy of the particles is sufficient to overcome the force of attraction between them, the particles in a gas are disordered and widely separated. The much greater separation of particles in a gas, compared to solids and liquids, can be seen when 1 mole of water (approximately 18cm3 at room temperature) gives 1 mole of steam (approx. 33000cm3 at 100oC and 1 atm.)

The particles in a gas are in constant and rapid motion in straight lines. They are constantly colliding with each other and with the walls of the container.

The pressure exerted by a gas depends on the number and the energy of the collisions which the particles make with the container walls per unit area per unit time. As the temperature of a gas increases, the mean kinetic energy of its particles increases. Collisions with the container walls are more energetic and occur more frequently, therefore the pressure increases.


Vapourpressure

Temperature

As the temperature of a liquid increases, the mean kinetic energy of its particles increases, and so more particles are able to escape to form a vapour. Therefore, as the temperature of a liquid increases, its vapour pressure increases.When the vapour pressure of the liquid becomes equal to the atmospheric pressure above it, the liquid boils.

A gas is confined only by its container. Therefore, it has no definite shape and no definite volume, taking up instead the shape and volume of its container.


Glossary of TermsWord / Term Definition / DescriptionIONIC BONDINGionic bond the attraction between oppositely charged ions formed

by electron transfer from a metal to a non-metalpolarising a description given to a cation of high charge density

(small and highly charged) e.g. Al3+ that can distort the spherical nature of an anion, causing an ionic bond to have covalent character

polarisable a description given to a large anion e.g. I- that has low control over its electron cloud due to high shielding of the nuclear charge, and which can be distorted away from its spherical shape by a small, highly charged cation

cation a positively charged ionanion a negatively charged ionCOVALENT BONDINGmolecule a group of two or more atoms held together by covalent

bonds e.g. Cl2, H2O, NH3, CH4, I2, BeCl2, BF3, PF5, SF6, AlCl3covalent bond a shared electron pair between two atomsdative covalent bond

a shared electron pair between two atoms where both electrons have originated from one atom (also known as a coordinate bond)

coordinate bond

a shared electron pair between two atoms where both electrons have originated from one atom (also known as a dative covalent bond).

electronegativity

the relative tendency of an atom to attract the electrons in a covalent bond

polar bond a covalent bond where the electron pair is shared unequally due to the difference in electronegativities of the atoms in the bond.

Lewis acid an electron pair acceptorLewis base an electron pair donorlone pair a pair of non-bonding electronsdimer a species formed between two small molecules e.g. Al2Cl6electron rich an atom which has a greater share of electron density

(shown by a - sign)electron deficient

an atom in a polar covalent bond which has the smallest share of the electron pair (shown by a + sign)

ionic character A covalent compound which shows some properties characteristic of an ionic compound. Caused by the unequal sharing of electrons in a covalent bond. Shown by polar covalent compounds.

polar molecule a polar molecule is a non-symmetrical molecule which contains polar bonds. If the molecule contains polar bonds but is symmetrical, the overall molecule is non-polar.

STRUCTURES OF SOLIDS CONTAINING COVALENT BONDSgiant covalent a type of structure. A lattice of atoms held together by


many strong covalent bonds. E.g. diamond, graphite, silicon, silicon dioxide.

simple molecular

a type of structure. A lattice of molecules held together by weak intermolecular forces.

intermolecular forces

the weak forces between molecules in a simple molecular solid.

van der Waals forces

the weakest intermolecular forces. A temporary, induced dipole on all molecules caused by temporary repulsions of the electron cloud in the molecules, causing the molecules to be weakly attracted to each other. The only type of intermolecular force present in solid iodine.

permanent dipole-dipole attractions

an intermolecular force. Weaker than hydrogen bonding, but stronger than van der Waals forces. The weak electrostatic attraction between the permanent dipoles on polar molecules.

hydrogen bonding

the strongest intermolecular force. The attraction between a very electron deficient hydrogen atom (which is covalently bonded to N, O or F) and the lone pair of electrons on a neighbouring molecule. Responsible for the high melting and boiling points of water.

ice solid water! A simple molecular solid held together by hydrogen bonding intermolecular forces.

SHAPES OF MOLECULESlinear the basic shape of a molecule with two electron pairs

around the central atom. Bond angle 180o. E.g. BeCl2trigonal planar the basic shape of a molecule with three electron pairs

around the central atom. Bond angles 120o. E.g. BCl3tetrahedral the basic shape of a molecule with four electron pairs

around the central atom. Bond angles 109.5o. E.g. CH4 trigonal bipyramidal

the basic shape of a molecule with five electron pairs around the central atom. Bond angles 120o and 90o. E.g. PCl5

octahedral the basic shape of a molecule with six electron pairs around the central atom. Bond angles 90o. E.g. SF6

pyramidal the shape of a molecule with 3 bonding pairs and one lone pair e.g. ammonia. The bond angle is reduced from the tetrahedral angle of 109.5o to 107o due to the greater repulsion of the lone pair of electrons.

bent planar the shape of a molecule with 2 bonding pairs and 2 lone pairs e.g. water. The bond angle is reduced from the tetrahedral angle of 109.5o to 104.5o due to the greater repulsion of the 2 lone pairs of electrons.

METALLIC BONDINGmetallic bond the electrostatic force of attraction between the regular

array of metal cations and the sea of delocalised outer electrons.

MISCELLANEOUS TERMScrystal a solid made up of a regular array (lattice) of particleslattice a regular pattern or 3-dimensional array of inter-linked

particles. A lattice can be used to describe the regular arrangement in any of the four main structure types.


Lattices can be ionic e.g. an alternating face centred cubic arrangement of Na+ and Cl- ions in NaCl held by electrostatic forces, giant covalent e.g. the regular tetrahedral arrangement of carbon atom in diamond held together by many strong covalent bonds, simple molecular e.g. the lattice of I2 molecules held together by weak van der Waals forces or metallic.

solid the lowest energy state of matter. Particles vibrate about a fixed position. On heating, the particles vibrate more vigorously until they have enough energy to move away from the fixed position and move around. Incompressible due to small spaces between particles.

liquid particles move randomly throughout the bulk of the liquid. There are still some forces of attraction acting between particles. Slightly compressible because particles have larger spaces between them.

gas particles move with rapid random motion. Easily compressible due to the large spaces between particles.

melting The change of state from a solid to a liquid.freezing The change of state from a liquid to a solid.boiling The change of state from a liquid to a gas.condensing The change of state from a gas to a liquid.sublimation The change of state from a solid straight to a gas

(missing out the liquid state). Can also be used to describe the change from a gas to a solid e.g. I2 (g) = I2 (s)

evaporation The change in state from a liquid to a gas below the boiling point of the substance.

PERIODICITY

CLASSIFICATION OF ELEMENTS

PERIODIC TRENDS IN PERIOD 3 (Na-Ar)


f-block

d-blockp-blocks-

block

1s

4f

3d

2p

1. Atomic Radius

the nuclear charge is increasing electrons are entering the same shell, so the shielding is constant the outer electrons are more strongly attracted and are drawn closer to the

nucleus


Na SMg ArClPSiAl

Atomic radius /nm

0.16

0.14

0.12

0.10

x

x

x

xx

xx

Atomic radius decreases across the period.

2. First Ionisation Energy

This trend has already been discussed in STRUCTURE AND BONDING.

3. Melting & Boiling Points

The melting and boiling points reflect the structure and bond strength of the elements.

Sodium, magnesium and aluminium are metals and have metallic bonding. From Na to Al, the size of the ion decreases due to the increasing nuclear charge, and the number of delocalised electrons increases. Therefore, the strength of the metallic bond increases. This results in an increase in both the melting point and the boiling point.

Silicon has a macromolecular structure in which there are strong covalent bonds in three dimensions, like diamond. Since a large amount of energy is required to break these bonds, the melting and boiling points are high.


Na SMg ArClPSiAl

First ionisation energy /kJ.mol-1

1600

1200

800

400x

x

x

x

xx

x

x

Temperature /K

Na SMg ArClPSiAl

3000

2000

1000

0

x

x x

x

x xx

2500

1500

500

x

xx

x x

xx

x x

xx

Tb

Tm

Phosphorus (P4), sulphur (S8) and chlorine (Cl2) are simple molecular substances with weak van der Waals’ forces holding the molecules together. Since van der Waals’ forces are weak, all three elements have fairly low melting and boiling points. The strength of the van der Waals’ forces increases as the size of the molecule increases. Therefore, sulphur, which has the largest and most polarisable molecule, has the highest melting and boiling points. This is followed by phosphorus, then by chlorine.

Argon is monatomic and has very weak van der Waals’ forces between its atoms. Its melting point (84K) and boiling point (87K) are both very low.


Documents

BONDING - Barnard Castle School€¦ · Web viewHydrogen bonding explains the unusually high melting and boiling points of molecules such as water. With an MR of only 18, water