Bonding and Structure

1. Bonding and Structure

2. How atoms bond together The bonds between atoms always involve their outer electrons • Inert gases have ____ outer shells of electrons and are very__________. • When atoms bond together they share, pool or transfer electrons to get full outer shells, making them more ______and less ________, like the inert gases. • There are three types of strong chemical bonds: ionic, covalent and ________. • Mostly we show these bonds through drawing _____ structures, these are ‘dot and cross diagrams’. metallic full reactive Lewis stable unreactive

3. Ionic Bonding

4. Contents Describe the ionic bond as the electrostatic attraction between oppositely charged ions. Describe how ions can be formed as a result of electron transfer. Deduce which ions will be formed when elements in group 1, 2 and 3 lose electrons. Deduce which ions will be formed when elements in groups 5, 6 and 7 gain electrons. State that transition elements can form more than one ion.

5. Ionic Bonding • Ionic bonding occurs between metals and non-metals and involves transfer of _________ from metals to non-metals. • Sodium chloride has ionic bonding. • Remember sodium has ___ electrons, its electron arrangement is _______. • Chlorine has _____ electrons, its electron arrangement is _________. • Now draw these atoms showing all the electrons. 17 electrons 2,8,1 11 2,8,7

6. An electron is transferred. The single outer electron of the sodium atom moves into the outer shell of the chlorine atom + - Na Cl + - 2, 8 , 1 2, 8, 8 7

7. + - Cl Na + - 2, 8 2, 8, 8 Each outer shell is now ____. Both sodium and chlorine have an inert gas electron structure. The two charged particles are called _____. The sodium ion is ________charged, because it has _____an electron. The chloride ion is ________charged, because it has _____an electron. The two ions are attracted to each other and to other oppositely charged ions in the sodium chloride compound by ____________forces gained full electrostatic positively lost ions negatively

8. The number of electrons lost or gained will depend on the number of electrons in the outer shell of the atom. 2- 2+ O Mg Mg has 2 electrons in its outer shell, it can therefore lose 2 electrons and become a 2+ ion Oxygen has 6 electrons in its outer shell, it can therefore gain 2 electrons and become a 2- ion

9. + 2- Na O + Look what happens when sodium bonds with oxygen? Sodium can only donate one electron, but oxygen needs two One oxygenatom needs to react with two sodium atoms to balance the charges. Na

10. In summary • Ionic bonding is the result of electrostatic attraction between oppositely charged ions. • The attraction extends throughout the compound making a structure called a lattice. • The formula for sodium chloride is NaCl, because we know that for every one sodium ion there is one chloride ion. • The formula for magnesium oxide is MgO, because there is one magnesium for every oxide. • The formula for sodium oxide however is Na2O because two sodium ions are needed for each oxide ion

11. Common ions formed Metals always form positive ions, by losing electrons Non-metals always form negative ions by gaining electrons +1 +3 +2 -1 -3 - 2 variable

12. Common ions formed in the transition elements

13. Contents State the formula of common polyatomic ions formed by non metals in period 2 and 3. Describe the lattice structure of ionic compounds.

14. Other polyatomic complex ions NO3- MnO4- OH- SO42- Cr2O72- PO43- CO32-

15. Properties of ionic compounds – melting point • Ionic compounds are always solids. They have giant structures and therefore high melting points. • This is because in order for them to melt an ionic compound, energy must be supplied to break up the lattice of ions.

16. SOLID - + - + + + + + - + - - - - - - + - Cathode - Anode + + - + - Properties of ionic compounds – conducting electricity • Ionic compounds conduct electricity when molten or dissolved in water (aqueous) but not when solid. • This is because the ions that carry the current are free to move in the liquid state but are not free in the solid state. LIQUID Anode + Cathode -

17. Properties of ionic compounds – hardness • Ionic compounds tend to be brittle and shatter easily when given a sharp blow. This is because they form a lattice of alternating positive and negative ions. • A blow in the direction shown may move the ions and produce contact between ions with like charges.

18. Covalent Bonding

19. Contents Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei. Describe how the covalent bond is formed as a result of electron sharing. Deduce the Lewis (dot and cross) structures of molecules and ions for up to four electron pairs on each atom. State and explain the relationship between the number of bonds, bond length and bond strength.

20. Covalent bonding • Covalent bonds form between __________atoms. • The atoms _____ some of their outer electrons so that each has a full outer shell of electrons. • A covalent bond is a shared _____ of electrons. • We can represent one pair of shared electrons in a covalent bond by a line, Cl – Cl. • Many non-metal elements exist as ________ molecules, covalently bonded to each other diatomic non-metal pair share

21. Forming molecules • A small group of covalently bonded atoms is called a molecule. • For example chlorine is a gas which is made of molecules, chlorine has 17 electrons and an electron arrangement 2,8,7. • Two chlorine atoms make a molecule. • The two atoms share one pair of electrons • Each atom now has a full outer shell. • The molecule does not contain charged particles because no electrons have been transferred from one atom to another.

22. Cl Cl Cl Cl The Cl – Cl bond is formed from the sharing of two electrons; one from each of the chlorine atoms to form the diatomic, covalent molecule, Cl2

23. Methane • Methane gas is a covalent compound of carbon and hydrogen. • Carbon has 6 electrons with electron arrangement 2,4 and hydrogen has just one electron. • In order for carbon to get a full outer shell, there are four hydrogen atoms to every carbon atom.

24. H Hydrogen has just 1 electron in its outer shell. H C H H Carbon has the electron arrangement 2,4 Four hydrogen's are needed for each carbon atom H

25. Double Covalent Bonds • In a double bond, four electrons are shared. • For example, the two atoms in an oxygen molecule share two pairs of electrons, so that the oxygen atoms have a double bond between them. O O O

26. Activity: • Draw the triple bond between two nitrogen atoms

27. Properties of simple covalent molecules • Substances composed of molecules are gases, liquids or solids with low melting points. (The strong covalent bonds are only between the atoms in the molecules, not between the molecules themselves). • They are poor conductors of electricity (because there are no charged particles to carry the current). • If they dissolve in water, the solutions are poor conductors of electricity. (Again there are no charged particles).

28. Dative Covalent Bonding • A single covalent bond consists of a pair of electrons shared between two atoms. • In most covalent bonds, each atom provides one of the electrons, but in some bonds, one atom provides both the electrons. • This is called dative covalent bonding, it is also called coordinate bonding. • In a dative bond; • The atom that receives the electrons is electron deficient. • The atom that is donating the electrons has a pair of electrons that is not being used in a bond, called a lone pair.

29. The ammonium ion • For example, ammonia, NH3, has a lone pair of electrons. Draw the Lewis structure of ammonia. Ammonium ion Lone pair H + H N H H+ H N H H H In the ammonium ion, NH4+, the nitrogen uses its lone pair of electrons to form a dative bond to an H+ ion (a ‘bare’ proton with no electrons at all and therefore electron deficient).

30. Dative covalent bonds are represented by an arrow, . The arrow points towards the atom that is receiving the electron pair. But this is only to show how the bond was made. • The ammonium ion is completely symmetrical and all the bonds have exactly the same strength and length. • Dative bonds have exactly the same strength and length as ordinary covalent bonds between the same pair of atoms. • The ammonium ion has covalently bonded atoms, but is a charged particle. Ions like this are called complex ions

31. Activity: • Draw Lewis structures showing all electrons for the following covalent compounds: • Water, H2O • Carbon dioxide, CO2 • Fluorine, F2 • Ethane, C2H6 • Ammonia, NH3

32. Shapes of covalent molecules

33. Contents Predict the shape and bond angles for species with four, three and two negative charge centers on the central atom using the valence shell electron pair repulsion theory (VSEPR). Predict whether or not a molecule is polar from its molecular shape and bond polarities.

34. The shapes of covalent molecules • Molecules are three dimensional and they come in many different shapes. • We can predict the shape of a simple covalent molecule – for example, one consisting of a central atom surrounded by a number of other atoms – by using the ideas that: • A group of electrons around an atom will repel all other electron groups. • The groups of electrons will therefore take up positions as far apart as possible. • This is called the Valence Shell Electron Pair Repulsion theory or VSEPR (the valence shell is another name for outer shell).

35. A group of electrons may be: • A set of two shared in a single bond • A set of four in a double bond • An unshared (lone) pair • The shape of a molecule depends on the number of groups of electrons that surround the central atoms. • To work out the shape of any molecule you must first draw a Lewis, ‘dot and cross’, diagram.

36. O O C Two groups of electrons • If there are two groups of electrons around the atom, the molecule will be linear. • The furthest away from each other the two groups can get from each other is 180o. • Carbon dioxide has this shape, because it has two groups of electrons. Each group has four electrons in a double bond between carbon and oxygen. 180o O = C = O

37. F 120o B F F Three groups of electrons • If there are three groups of electrons around the central atom, they will be 120oapart. • The molecule is flat and the shape is trigonal planar. Boron trifluoride is an example of this: F B F F Outer electrons shown only

38. H 109.5o C H H H Four groups of electrons • If there are four groups of electrons, they are furthest apart when they are arranged so that they point to the four corners of a tetrahedron (a triangular based pyramid) The angles here are 109.5o. • Methane, CH4, is an example: H H H C H Outer electrons shown only

39. + H 109.5o N H H H Ammonium ion • The ammonium ion is also shaped like a tetrahedron. • It has four groups of electrons surrounding the nitrogen atom. + H H N H Outer electrons shown only H

40. Molecules with lone pairs • Some molecules have unshared (lone) pairs of electrons. These are electrons that are not part of a bond. • The lone pairs affect the shape of the molecule. • Water and ammonia are good examples of this effect. • There is an increase in the repulsion between the following groups: • Bonding pair – bonding pair • Lone pair – bonding pair • Lone pair – lone pair Repulsion increases

41. H N H H Ammonia, NH3 • Ammonia has four groups of electrons and one of the groups is a lone pair. With four groups, the ammonia molecule, like the water molecule, has a shape based on a tetrahedron. In this case there are only three ‘arms’. The lone pair squeezes the three shared pairs together and the bond angles are approximately 107o. The shape is described as pyramidal Outer electrons shown only N H H H 107o

42. Water, H2O H O • Look at the Lewis diagram . There are four groups of electrons around the oxygen so the shape is based on a tetrahedron. • However, two of the ‘arms’ of the tetrahedron are lone pairs that are not part of a bond. • This results in a ‘V’ shaped or angular molecule. O H H Outer electrons shown only H 104.5o

43. Lone pairs explained • The angles of a perfect tetrahedron are all 109.5o but lone pairs affect these angles. • The shared pairs of electrons are attracted towards the oxygen nucleus and also the hydrogen nucleus. • However, lone pairs are attracted only by the oxygen nucleus and are therefore pulled closer to it than shared pairs. • The lone pairs therefore repel more effectively than shared pairs, and ‘squeeze’ the hydrogen together, reducing the H-O-H angle. • An approximate rule of thumb is 2o per lone pair. • The actual bond angle in water is about 104.5o.

44. AHL Shapes of molecules and ions

45. Contents • Predict the shape and bond angles for species with five and six negative charge centres using the VSPER theory.

46. Shapes of five and six negative charge centres • VSEPR theory can be extended to cover five and six pairs of electrons or negative charge centres. • The electron pairs will arrange themselves around the central atom so that they are as mutually repulsive as possible. • Five pairs will give a trigonal bipyramid shape with bond angles of 90˚, 120 ˚, 180˚ e.g. phosphorus pentachloride PCl5 • Six pairs will give an octahedral shape with bond angles of 90˚ and 180˚ e.g. sulphur hexafluoride SF6

47. Activity • Use the orbit molecular building system to create a trigonal bipyramid shape and an octahedral shape. F Cl Cl F F S P Cl Cl F F Cl F Trigonal bypyramid, PCl5 Octahedral, SF6

48. Expanding the Octet • The presence of five or six pairs of electrons around the central atom implies that the octet has been expanded. This cannot happen with the second period elements such as nitrogen, oxygen or fluorine, but can happen with the third period elements such as phosphorus, sulphur and chlorine. • This is because, in the third period, the elements have 3d orbitals, which are close enough in energy to the 3p orbitals that they are available to be utilised. • Nitrogen, for example, forms only one chloride, NCl3, whereas phosphorus can form two chlorides, PCl3 and PCl5.

49. Non-bonding pairs • When there are non-bonding pairs of electrons then the same rules apply, in that non-bonding pairs exert a greater repulsion than bonding pairs. • Consider the structure of xenon tetrafluoride (XeF4). • There are six pairs of electrons around the central xenon atom. Four of them are bonding pairs and two are non-bonding pairs. • This could result in two possible structures. In the first structure the bond angle between the non-bonding pairs is 90˚, whereas in the second structure it is 180˚, attempt to draw these two structures. Which structure will exist? What is the name of this shape? • The actual shape of xenon tetrafluoride is therefore square planar

50. Bond Length, Strength and Resonance Hybrids