West Midlands Chemistry Teachers Centre November 2009 Structure and Bonding Presenter: Dr Janice Perkins

Ionic Bonding Covalent Bonding Metallic Bonding Types of Bonding Ionic Bonding Covalent Bonding Metallic Bonding

Ionic Bonding Negative Ions (anions) have a negative charge because of a surplus of e- Positive Ions (cations) have a positive charge because of a deficiency of e- Ions formed by e- transfer from one atom to another the number of e- lost or gained depends on the elements involved Ionic Bonding is the attraction between these positive and negative ions It is called ELECTROSTATIC ATTRACTION

Question from past paper Magnesium and chlorine react together to form the ionic compound magnesium chloride, MgCl2 (i) Explain how each of the ions in this compound is formed (ii) Explain why compounds with ionic bonding tend to have high melting points Two e- transferred from Mg to Cl  one e- to each of two Cl atoms   electrostatic attractions are strong  Too easy? No, it was very poorly answered

Covalent Bond Covalent bond = shared pair of e- One electron comes from each atom

Question from June 09 Two organic compounds with similar relative molecular masses are shown below. Ethanol Propane State the type of bond present between the C and H atoms in both of these molecules. Explain how this type of bond is formed. Covalent Shared pair of electrons, one from each atom 

also called Dative covalent bonding Co-ordinate Bonding also called Dative covalent bonding Co-ordinate bonding is to do with how a covalent bond is formed Once formed it is a normal covalent bond One atom/ion supplies both e- to another atom/ion

Question from Jan 09 Phosphorus is in the same group of the Periodic Table as nitrogen. The molecule PH3 reacts with an H+ ion to form a PH4+ ion. Name of the type of bond formed when PH3 reacts with H+ and explain how this bond is formed. Coordinate/dative  Both electrons/ lone pair (on P/PH3)  Shares/donated from P(H3)/ to H(+) 

Metallic bonding This type of bonding is often very poorly explained. frequently see: ionic bonds hydrogen bonds van der Waals’ forces in the answer of even the better candidates This is wrong!

Metallic bonding + + + + + + + + X X + + + + X + + + + X OUTER electrons are ‘DELOCALISED’ i.e. free to move through metal Metal lattice held together by the ATTRACTION between this sea of delocalised e- and the positive ions in the lattice

Question from June 09 State the type of bonding involved in silver. Draw a diagram to show how the particles are arranged in a silver lattice and show the charges on the particles. metallic bonding  regular arrangement of same sized particles  + charge in each ion 

Bond Polarity Cl x Cl A bond is formed between atoms of the same element They both attract the electrons in the bond to the same extent This attraction is called electronegativity

Electronegativity Definition: Trends: The power of an atom to attract electron density (or a pair of electrons)  from a covalent bond.  Trends: Electronegativity increases across the period Due to more protons and smaller size, so stronger attraction for the bonding e- Decreases down a group Due to larger size/more shells, so weaker attraction for bonding e-

What’s the sporting Caption? ‘Throwing the puppy dog!’??

What’s the sporting Caption? Wish we were on holiday

What’s the sporting Caption? Well this one is clear enough It must be ‘Tug-o-War’

Electronegativity and covalent bonds Like a tug-o-war The ‘tug’ is provided by the electronegativity of the atom A tug-o-war between equal teams is like the pull of the Cl atoms in the covalent bond because atoms of the same element have same electronegativity. perfectly even sharing of the bonding e- Cl x Cl

Covalent bonds between atoms with different electronegativities The attraction for the bonding e- by the two atoms is different This results in unequal sharing of the bonding e- between the two atoms The extent of the unequal distribution will depend on how different the electronegativities are - the term we use is ‘electronegativity difference’ The bigger the difference, the greater the dipole

AB is a polar molecule it has a dipole Polar Bonds Atom B is more electronegative than atom A Bonding e- are attracted more strongly to B + - Electron density is low. Electron density is high. A x B +A – B - AB is a polar molecule it has a dipole

Forces acting between molecules Types of Intermolecular Forces (IMF) van der Waals’ forces (temporary induced dipole - dipole attractions) Permanent dipole - dipole attractions Hydrogen bonding All result from attractions between the partial charges in the dipoles The stronger the dipole – the stronger the IMF

Van der Waals’ forces It is the weakest of the three IMFs It result from temporary unequal distributions of e- density The bigger the atom or molecule, the more electrons there will be and the larger will be the surface area – this increases the strength of the van der Waals’ attractions It is always present but is often swamped by stronger IMFs It is the only IMF present between non-polar molecules

Permanent dipole-dipole forces Always refer to them as ‘dipole-dipole’, not just ‘dipole’, attractions They result from attractions between the and of polar molecules The greater the electronegativity difference between the two atoms in the bond, the greater the dipole + -

Hydrogen Bonding This is the strongest IMF. It is an extreme example of dipole-dipole attractions It only occurs in molecules with very large electronegativity differences. N–H, O–H and F–H bonds It results in an attraction that is about 10% as strong as a covalent bond

Hydrogen bonding in HF δ+ H F δ- δ+ H F δ- δ+ H F δ- H F H F  Mark 1 = 3 lone pairs on the fluorine atoms Mark 2 = dipole correctly shown   Mark 3 = hydrogen bond between lp and H δ+ H F δ- δ+ H F δ- δ+ H F δ- H F H F

Question from Jan 09 H C N O Electronegativity 2.1 2.5 3.0 3.5 State the strongest type of intermolecular force in the following compounds. Methane (CH4) van der Waals Ammonia (NH3) Hydrogen bonding Use the values in the table to explain how the strongest type of intermolecular force arises between two molecules of ammonia Large electronegativity difference between N + H  Forms N - / H +  Lone pair on N attracts H (+) 

States of Matter The properties shown by materials are the result of their structure and bonding You need to know about crystalline materials which are: Ionic Metallic Giant covalent Molecular

Ionic Crystal Lattices Each ion is surrounded by a number of ions of the opposite charge In sodium chloride, each ion is surrounded by 6 of the oppositely charged ions. This gives a cubic shape. Remember. Oppositely charged ions attract each other by Electrostatic attraction

2D diagram of NaCl Each Na+ ion is surrounded by a total of 6 Cl- ions One more is be in front of the ion; the other is behind it

3D diagram of NaCl Cl Cl- - Each Na+ ion is surrounded by six Cl- ions and vice versa

Why are ionic solids brittle animation?

Ionic Lattice OK – so what was happening there? Here it is again with sub-captions for the hard-of-thinking!

Why are ionic solids brittle animation? Like charges repel

Metals In Mg there is an attraction between the delocalised e and the lattice of metal ions 2+ 2+ 2 + 2 + 2+ 2+ 2 + 2 + 2+ 2+ 2 + 2 + Most metals have high m.p. / b.p. as a lot of energy is needed to remove an atom from this attraction

The delocalised e- flow thorough the metal (a current) so metals conduct electricity

So how can metals change their shape? Malleable since the layers can slide over one another

Comparison of Na, Mg and Al delocalised e- Na+ Mg2+ Al3+ All ions have the same electron configuration (isoelectronic) As the proton number increases from Na to Al, so does the attraction for the outer shell electrons. Size decreases Higher ionic charge, more delocalised e-, smaller ionic radius Metallic bonding increases; so m.p. / b.p. increase Na  Al

Giant molecular crystals Diamond Graphite Silica

Physical strength explain in terms of bond breaking (diamond/Si) and layers sliding (graphite) Many strong covalent bonds need to be broken in the structure of diamond and silica so strong substances. In graphite the layers can slide over one another so graphite is softer and is often used as a lubricant. Melting points explain in terms of bond breaking Many strong covalent bonds need to be broken in macromolecules so high melting points. Electrical conductivity explain in terms of presence/absence of delocalised e- Graphite has delocalised electrons between the layers so it conducts electricity.

Molecular crystals The example you need to know is Iodine (but you could be asked to apply your understanding to other elements, such as sulphur) There is only one element present The bond is non-polar The only IMF present is van Der Waals’ These IMFs are weak so The melting point is relatively low The IMFs operate between molecules

Shapes of molecules/ions The basic shape is determined by the number of electron pairs present If all the electron pairs are bonding-pairs then we get a ‘regular’ shape If some of the electron pairs are lone-pairs then we get an ‘irregular’ shape

The Electron Pair Repulsion Theory The shape results from repulsions between e- pairs – NOT bonds/atoms The rules are: Lone pairs repel more strongly than bonding pairs Order is l.p/l.p. > l.p/b.p >b.p/b.p Remember: When lone pairs are present, the basic shape will be distorted.

Equal repulsion between Linear Shape B A B Equal repulsion between 2 bonding pairs

Equal repulsion between Trigonal Planar shape Equal repulsion between 3 bonding pairs

Equal repulsion between Tetrahedral shape B A Equal repulsion between 4 bonding pairs

Trigonal Bipyramid shape Equal repulsion between 5 bonding pairs

Equal repulsion between Octahedral shape Equal repulsion between 6 bonding pairs

Regular Shapes 2 B – A – B 3 AB3 4 AB4 5 AB5 6 AB6 180 linear 120 e- pairs Formula Bond angle Name 2 B – A – B 180 linear 3 AB3 120 trigonal planar 4 AB4 109½ tetrahedral 5 AB5 90 trigonal bipyramid 6 AB6 octahedral

Working out the shape If it is an ion You can work out the shape using simple maths Find the group number of the central atom Add to this number the number of bonds present If it is an ion add one if the charge on the ion is -1 Deduct 1 if the charge is +1 Divide total by 2 - this gives the number of e- pairs Number of e- pairs = basic shape If lone pairs present, basic shape modified

To draw the shape Work out the basic shape Lightly sketch this in pencil Put the atoms in Draw the lone pairs along the ‘spare’ bonds Rub out you working pencil lines

Working out the shape of NH3 N is in Group 5, so five outer electrons 3 more electrons (1 from each H) Makes 8 electrons 4 pairs of electrons 3 bonding pairs and 1 lone pair Basic shape (4 e- pairs) = tetrahedral One lone pair, so doesn’t have basic shape 3 bonding pairs + 1 lone pair

Ammonia Pyramidal shape Unequal repulsion between 3 bonding pairs and 1 lone pair Pyramidal shape

Working out the shape of the ion NH2- N is in Group 5, so five outer electrons 2 bonds to H atoms, so 2 extra bonding e- 1 charge on molecule, so add extra e- = 8 electrons Makes 2 bonding pairs and 2 lone pairs Basic shape (4 e- pairs) = tetrahedral 2 bonding pairs + 2 lone pair so doesn’t have basic shape

The shape of the ion NH2- ‘Bent’ or ‘V-shaped’ Unequal repulsion between 2 lone pair and 2 bonding pairs ‘Bent’ or ‘V-shaped’

Shapes of molecules / ions

Question from Jan 09 Arsenic is in the same group as nitrogen. It forms a compound AsH3 Draw the shape of an AsH3 molecule and name the shape made by its atoms. 3 bonds and 1 lp attached to As trigonal pyramidal  As H ·

The End