Bonding HL and SL

4.1 Ionic Bonding An ionic bond is an electrostatic attraction between ions of opposite charge. Positive ions (or cations) are formed when electrons are lost from the outer energy level of, usually, a metal atom. Negative ions (or anions) are formed when electrons are gained to the outer energy level of a non-metal atom. Ions usually have filled outer energy levels (except certain transition metal ions). Transition metals can form more than one positive ions e.g. Fe 2+ and Fe 3+, Cu + and Cu 2+.

Positive ions are formed when atoms lose electrons. e.g. Li(2, 1)  Li + (2) + e - Li(1s 2 2s 1 )  Li + (1s 2 ) + e - HL only Do similar to show the formation of the following ions: Na + Mg 2+ K + Ca 2+ Al 3+ Cu 2+ and Fe 3+ (HL only) Think carefully about the number of electrons that need to be lost. Is there a link between the formula of the ion and the position of the atom in the periodic table?

Negative ions are formed when atoms gain electrons. e.g. F(2, 7) + e -  F - (2, 8) F(1s 2 2s 2 2p 5 ) + e -  F - (1s 2 2s 2 2p 6 )HL only Do similar to show the formation of the following ions: Cl - O 2 - S 2 - N 3 - P 3 - Br - (HL only) Think carefully about the number of electrons that need to be gained. Is there a link between the ion formed and the position of the element in the periodic table?

All of the ions considered so far are made up of just one atom but we should know the formula of the following polyatomic ions: NO 3 - nitrate(V) ion OH - hydroxide ion SO 4 2- sulfate(VI) ion CO 3 2- carbonate ion PO 4 3- phosphate(V) ion HCO 3 - hydrogencarbonate ion NH 4 + ammonium ion Brackets are used in formulae to show more than 1!

Ionic compounds are formed when electrons are transferred from a metal atom to a non-metal atom. A compound will be ionic if the difference in electronegativity of the two atoms is greater than 1.8 (see later). Look at the model for the formation of sodium chloride on the next slide. What is wrong with this model?

a sodium atom (Na) 2, 8, 1 17 p 18 n a chlorine atom (Cl) 2, 8, 7 11 p 12 n chemical reaction + 11 p 12 n a sodium ion (Na + ) (2, 8) + a chloride ion (Cl - ) (2, 8, 8) - 17 p 18 n - The formula of sodium chloride is NaCl

12 p 12 n a magnesium atom (Mg) 2, 8, 2 8 p 8 n an oxygen atom (O) 2, 6 chemical reaction p 12 n a magnesium ion (Mg 2+ ) (2, 8) 2+ 8 p 8 n 2- a oxide ion (O 2- ) (2, 8) 2- The formula of magnesium oxide is MgO

calcium chloride 17 p 18 n a chlorine atom (Cl) 2, 8, 7 a calcium atom (Ca) 2, 8, 8, 2 20 p 20 n The calcium atom has two electrons in its outer shell which must both be lost for it to have an electronic structure of a noble gas. The chlorine atom only has space for one more electron in its outer shell. What happens?

The calcium atom reacts with 2 chlorine atoms 20 p 20 n 17 p 18 n 17 p 18 n

17 p 18 n - 17 p 18 n p 20 n The formula of calcium chloride is CaCl 2

Compounds which contain ionic bonds are solids at room temperature. Ionic bonds do not exist in isolation but as part of a GIANT IONIC LATTICE. In a giant ionic lattice each positive ion is attracted by negative ions which surround the positive ion in a regular arrangement and vice versa. This attraction is an electrostatic attraction. e.g. sodium chloride

Or in 2-d: sodium ion chloride ion Ionic compounds conduct electricity when molten or in aqueous solution. They do not conduct when solid as the ions are held in fixed lattice positions. When liquid or in aqueous solution, the ions are free to move.

4.2 Covalent Bonding Covalent bonding is an alternative way for atoms to achieve a filled outer energy level. Covalent bonds are formed between non-metal atoms or atoms with an electronegativity difference of less than 1.8. A covalent bond is an electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the two atoms held by the bond. Often shown as H – F.

Here are some hints to help draw Lewis structures: 1.Work out the number of shared pairs required by each atom. This is the same as the number of extra electrons required to fill the outer shell. 2.Draw a diagram showing the bonds between atoms only e.g. Cl – Cl. 3.Draw an outline dot-cross diagram. 4.Fill in the electrons in the covalent bonds. 5.Make sure that the outer shells are filled. Other examples include: These diagrams are called LEWIS STRUCTURES.

Draw Lewis structures for: 1. Hydrogen chloride, HCl7. Carbon dioxide, CO 2 2. Oxygen, O 2 8. Hydrogen cyanide, HCN 3. Tetrachloromethane, CCl 4 9. Carbon disulphide, CS 2 4. Hydrogen sulphide, H 2 S10. Ethene, C 2 H 4 5. Phosphorus(III) chloride, PCl Ethyne, C 2 H 2 6. Nitrogen, N Carbon monoxide, CO

Co-ordinate bonding Aka dative covalency A co-ordinate bond is formed when both electrons in a shared pair originate from one atom. This pair is called a lone pair on the donor atom. Once a co-ordinate bond is formed it cannot be distinguished from a normal covalent bond. Co-ordinate bonds are found in transition metal complex ions (HL see later), the ammonium ion, and the hydroxonium ion.

The ammonium ion is formed when ammonia reacts with H + + lone pair of electrons H + has no electrons H+H+ H + ammonium ion co-ordinate bond

H 3 O + (the hydroxonium ion) is formed when water reacts with H +. Draw similar diagrams to show the formation of the hydroxonium ion. Can H 4 O 2+ be formed? In theory yes but in practice no! Why not? Other examples of substance containing co-ordinate bonds are carbon monoxide (CO) and aluminium chloride (Al 2 Cl 6 ). Draw Lewis structures for these compounds.

Bond Length and Bond Strength Tabulate the bond length and bond strength of the following bonds: O – OO = O N – NN = NN  N C – CC = CC  C What patterns do you notice? Find out the C to O bond lengths in either the ethanoate ion or the carbonate ion. How do you explain this? Try drawing Lewis structures for these substances to help.

Other substances for which resonance hybrid structures can be drawn include: Sulphur dioxide, SO 2 Ozone, O 3 Nitrate ion, NO 3 - Benzene, C 6 H 6 Can you draw Lewis structures for them?

Ionic or covalent? Generally: Ionic if electronegativity difference between atoms > 1.8 Covalent if electronegativity difference between atoms < 1.8 However, some compounds are described as ionic with some covalent character whilst others are described covalent with some ionic character. Most compounds have bonding somewhere between the two extreme models of ionic and covalent bonding. Look at the substances properties, if they are given, to help you decide. For example: if it conducts when aqueous or molten it must be ionic.

Polar Covalent bonds Arises due to differences in electronegativity between the two atoms held by a covalent bond. Electronegativity is defined as the power of an atom to withdraw electron density from a covalent bond. Small atoms with a large number of protons attract electron density most strongly. So electronegativity in the periodic table increases from left to right bottom to top

H 2.1 He Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Ne Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 Ar Br 2.8 Kr increases See later in Periodicity topic.

When a covalent bond exists between atoms of differing electronegativities, the shared pair is displaced towards the more electronegative atom. AA ‘normal’ covalent bond with electron pair midway between the two atoms AB B is more electronegative than A so the electron pair is displaced towards B. The result is a polar covalent bond.

The displacement of electron density makes the less electronegative atom slightly electron deficient so has a small positive charge shown as  +. The more electronegative atom has a slight excess of electron density which is shown as  -. AB ++ -- If the molecule is placed between electric plates the  - end is attracted to the positive plate and the  + end is attracted to the negative plate. The molecule is said to have a dipole moment, the bigger the difference in electronegativities, the bigger the dipole moment.

Examples of molecules with polar bonds include: HF ++ -- CCl ++ -- H3H3

Shapes of molecules: If a molecule contains polar bonds, it does not necessarily follow that the molecule itself is polar. If the molecule is symmetrical then the dipoles will cancel each other out. So we need to be able to deduce the shape of a molecule before we can decide whether or not it is a polar molecule. To do this we use the VALENCE SHELL ELECTRON PAIR REPULSION (VSPER) theory. For the purposes of this theory we must treat double and triple covalent bonds as if they are single bonds.

The outer electrons of atoms are arranged in pairs. These electron pairs can be considered as ‘clouds’ of electron density which repel each other as far apart as possible. Shapes of molecules depend upon number of electron pairs around the central atom in an ion or molecule. n.b. it is electron pairs that repel each other not the atoms.

2 pairs

3 pairs

4 pairs

Rules to help determine number of electron pairs around central atom: Add together: number of electrons in outer shell of central atom when not bonded, number of shared pairs of electrons (same as number of surrounding atoms) Then if it is an ion Add one for each negative charge, subtract one for each positive charge.

Some examples: BeCl 2 There are 2 electrons from the outer shell of beryllium + 1 from each of the chlorine atoms. Total 4 electrons i.e. 2 pairs of electrons Two electron pairs repel each other equally so shape is linear. Bond angle 180 °.

CH 4 There are 4 electrons from the outer shell of carbon + 1 from each of the hydrogen atoms. Total 8 electrons i.e. 4 pairs of electrons Four electron pairs repel each other equally so shape is tetrahedral. Bond angle °.

Lone pairs of electrons are shorter and fatter than pairs involved in bonding. Why? As a result they repel more than bonding pairs leading to a reduction in bond angles between bonding pairs. So in ammonia, NH 3 5 electrons from N + 3 from shared pairs = 8  4 pairs But only 3 involved in bonding so there is one lone pair which reduces the bond angle to 107 °. So ammonia is:

This shape is sometimes referred to as pyramidal. So in water, H 2 O 6 electrons from O + 2 from shared pairs = 8  4 pairs But only 2 involved in bonding so there are two lone pairs which reduces the bond angle to 105 °. So water is:

This shape is sometimes referred to as v-shaped or bent.

4.3 Intermolecular Forces Intermolecular forces are weak forces between molecules. There are 3 types ( in order of increasing strength): van der Waals forces, permanent dipole – permanent dipole attractions, hydrogen bonding.

Van der Waals forces Aka temporary dipole – induced dipole attractions. Found in all molecules. About 1 % strength of a covalent bond. Arises due to fluctuations of electron clouds. At any instant there may be more electron density on one part of the molecule (lhs on illustration)  small negative charge,  -, on part of the molecule and small positive charge on another part,  +. -- ++ Temporary dipole

This causes an adjacent molecule to have a dipole (an induced dipole) as the negative charge in the temporary dipole repels electron density in the neighbouring molecule. -- ++ -- ++ Temporary dipole Induced dipole Forces only act for a short time as the electron density is continually changing. These forces are continually switched on and off.

Consequences of Van der Waals forces: 1.Increase in boiling point of alkanes with M r. The larger the molecule, the bigger the magnitude of van der Waals forces. See later. 2.Variation in b.p. of isomers of pentane. See later 185 K 309 K 283 K

Dipole – Dipole attractions Molecules with permanent dipoles attract each other as shown below: H – Cl ++ -- Cl – H -- ++ H – Cl ++ -- = weak electrostatic attractions

Hydrogen Bonding The big daddy of them all but still only maximum of 1/10 th strength of a covalent bond. It is a special case of a dipole – dipole attraction force which exists between a lone pair of electrons on a N, O or F atom and a hydrogen atom with a strong partial positive charge (  +) as it is attached to an atom with a large electronegativity (N, O or F) The electronegative atom pulls electron density from the hydrogen atom to such degree that the hydrogen almost appears like an unshielded proton.

e.g in water hydrogen bond Note that the nucleus of the hydrogen bonded hydrogen atom is always in line with the nuclei of the two electronegative atoms on either side. 180 °

Hydrogen bonding is also found in: HF, hydrogen fluoride NH 3, ammonia C 2 H 5 OH, ethanol CH 3 COOH, ethanoic acid In each of these cases the molecules have relatively high boiling points for such small molecules.

What is the trend in boiling points of the hydrides of groups 4, 5, 6 and 7? i.e for group 6: H 2 O, H 2 S, H 2 Se, H 2 Te and, for group 7: HF, HCl, HBr, HI Boiling points should increase as molecules increase in size down the group so the van der Waals forces increase. But ………. Plot a graph of period number against boiling point for each of these groups on one set of axes.

H 2 O and HF have higher boiling points than expected due to hydrogen bonding

Why is hydrogen bonding essential to life?

4.2 Covalent Bonding (more!) Giant covalent structures Carbon exhibits allotropy (it exists in more than 1 form in the solid state). The allotropes of carbon are DIAMOND, GRAPHITE and the FULLERENES.

Diamond Graphite Bucky ball

In diamond each C atom is bonded covalently to 4 others in a giant tetrahedral structure. Diamond is extremely hard as a consequence of the strong covalent bonds within the molecule. To melt diamond requires a lot of energy as all of the covalent bonds have to be broken. There are no free electrons so diamond does not conduct electricity. Silicon and silicon dioxide also have a diamond structure

In graphite each C atom is bonded covalently to 3 others in a trigonal planar arrangement forming hexagonal rings in layers. The forces between the layers are weak as they are formed by delocalised electrons between the layers (the 4 th electron from the outer shell of C). Graphite conducts electricity due to these delocalised electrons. As the forces are weak between the layers graphite feels slippery. The layer can move over each other so graphite can be used as a lubricant. It also has a high melting point as there are lots of strong covalent bonds to break.

The fullerenes were discovered jointly by Harold Kroto and Richard Smalley in In 1997 they were awarded the Nobel Chemistry Prize. The basic fullerene, C60, is generally known as buckminster fullerene (or a bucky ball). It consists of 60 carbon atoms joined in a combination of hexagonal and pentagonal rings (like a football).

4.4 Metallic Bonding Metallic bonds do not exist in isolation. They form part of a giant metallic lattice. A giant metallic lattice consists of close-packed metal ions surrounded by delocalised electrons. These delocalised electrons are free to move through the lattice. Where do they come from?

Metallic Bonding Metallic bonds do not exist in isolation. They form part of a giant metallic lattice. A giant metallic lattice consists of close-packed metal ions surrounded by delocalised electrons. These delocalised electrons are free to move through the lattice. Where do they come from?

4.5 Physical Properties Type of bonding StructureExampleProperties CovalentSmall molecule with van der Waals forces CovalentSmall molecule with dipole – dipole forces CovalentSmall molecule with hydrogen bonding CovalentGiant lattice of atoms IonicGiant lattice of positive and negative ions MetallicGiant lattice of metal ions surrounded by delocalised electrons