Bonding

1. Bonding

2. Chemical bonding

3. Ionic bonding • Occurs when one or more electrons are transferred from the outer shell of one atom to the outer shell of another atom (both to gain stable duplet or octet structure). • Results from the electrostatic attraction between cations and anions.

4. Example Sodium atom Na 2,8,1 Chlorine atom Cl 2,8,7 Sodium ion Na+ 2,8 Chloride ion Cl- 2,8,8 Na  [Na]+ + e- (Lewis structure) Cl + e-  [ Cl ]- Na x + Cl  [Na]+[ Cl x]- Sodium metal reacts with chlorine gas in a violent exothermic reaction to produce NaCl.

5. Formation of ions What ions will be formed by these elements? Na , Mg , Al , P , S and Cl Draw the dot and cross diagram for magnesium oxide

6. Ions of transition elements • The transition elements form more than one stable positive ion.

7. Preicting the type of bonding from electronegativity values • Ionic bonding typically occurs between metal and non-metal. E.g. Barium fluoride, BaF2 • The reactivity of metals and non-metals can be assessed using electronegativity (ability of an atom in a covalent bond to attract shared pairs of electrons to itself).

8. Fluorine, which has the greatest attraction for electrons in bond-forming situations is assigned the highest value on this scale. All other atoms are assigned values less than that of fluorine as shown above. • Note the following trends: • Metals generally have low electronegativity values, while non-metals have relatively high electronegativity values. • Electronegativity values generally increase from left to right within the Periodic Table of the elements. • Electronegativity values generally decrease from top to within each family of elements within the Periodic Table.

9. If the difference in E values is > 1.8 => ionic bond • If the difference in E values is 0, non-polar covalent bond • If the difference is 0 – 1.8, polar covalent bond

10. Electrons are not shared. E.g. Na+ Cl- , electron is transferred. Polar covalent bonds are covalent bonds with ionic character. Ionic bond Electrons are equally shared. E.g.Cl-Cl Non polar covalent bond Electrons are not equally shared. E.g. Atoms have different electronegativity values Polar covalent bond

11. Example • Use the table above to predict the type of bonding beween Fluorine , F2 Hydrogen iodide, HI and Lithium fluoride, LiF

12. Polyatomic ions

13. Structure of giant ionic compound • In an ionic compound, constituent ions are held in fixed positions in an orderly arrangementt by strong ionic bonds. Lattice structure consisting of a regular array of positively and negatively harged ions.

14. Covalent bonding • Formed by equal sharing of electrons between non-metallic elements to achieve the stable electronic configuration of noble gases • the shared electrons are localised between the two nuclei the attraction between the localised shared electrons and the nuclei is known as a covalent bond. • In the Hydrogen molecule, a bond between two atoms is formed by the sharing of electrons between the atoms.

15. The bonding pair of electrons spends most of its time between the two atomic nuclei, thereby screening the positive charges from one another and enabling the nuclei to come closer together than if the bonding electrons were absent. Negative charge on the electron pair attracts both nuclei and holds them together in a covalent bond. • From an energy standpoint, when we say two atoms are chemically bonded we mean the two atoms close together have less energy and therefore are more stable than when separated. • Energy given off by the atoms form a bond, and energy must be supplied to pull them apart. • A covalent bond is the result of electrostatic attraction between the nuclei of the 2 atoms and the pair of shared electrons.

16. Lewis structures Draw Lewis structures of the following molecules: Chlorine, Cl2 Hydrogen chloride, HCl Methane, CH4 Oxygen, O2 Nitrogen, N2 Carbon dioxide, CO2 Water, H2O

19. Coordinate (dative) bonding • In some molecules and polyatomic ions, both electrons to be shared come from the same atom. The covalent formed is called the coordinate or dative bond. • Carbon monoxide (CO) can be viewed as containing one coordinate bond and two "normal" covalent bonds between the C atom and the O atom.

20. Reaction between ammonia and hydrogen chloride • A thick white smoke of solid ammonium chloride is formed in the reaction below: • Ammonium ions, NH4+, are formed by the transfer of a hydrogen ion from the hydrogen chloride to the lone pair of electrons on the ammonia molecule.

21. When the ammonium ion, NH4+, is formed, the fourth hydrogen is attached by a dative covalent bond, because only the hydrogen's nucleus is transferred from the chlorine to the nitrogen. The hydrogen's electron is left behind on the chlorine to form a negative chloride ion. • Once the ammonium ion has been formed it is impossible to tell any difference between the dative covalent and the ordinary covalent bonds.

22. Dissolving hydrogen chloride gas in water • Something similar happens. A hydrogen ion (H+) is transferred from the chlorine to one of the lone pairs on the oxygen atom. • The H3O+ ion is variously called the hydroxonium ion.

23. Other examples: • The reaction between ammonia and boron trifluoride, BF3

24. Draw lewis structures for molecules and ions Rules • Calculate the total no. of valence electrons for all atoms in the molecule or ion. • Arrange all the atoms surrounding the central atom by using a pair of electrons per bond.The central atom is usually the atom that is least electronegative. [not H] • Assign the remaining electrons to the terminal atoms so that each terminal atom has 8 electrons. [ H will only have 2 ] • Place any electrons left over on the central atom. [P and S elets from period 3 may have > 8 electrons] • Form multiple bonds if there are not enough electrons to give the central atom an octet of electrons.

25. Example Write the Lewis structure (electron dot diagram) for hydrogen cyanide, HCN

26. Bond strength and length of covalent bonds • Strength Triple bonds > Double bonds > Single bonds • Length Single bonds > Double bonds > Triple bonds

27. Bond polarity • In diatomic molecules (e.g. H2 ,Cl2) both atoms exert an identical attraction. • When the atoms are different (e.g. HCl) with one more electronegative than the other, a polar bond is formed. • Relative polarity is predicted from electronegativity values. • C-O is more polar than C-Cl since the difference in E value for C-O is greater than that for C-Cl.

28. VSEPR theory • The shapes of simple molecules and ions can be determined by using the Valence Shell Electron Repulsion (VSEPR) theory. • Electron pairs around the central atom repel each other • Bonding pairs and lone pairs arrange themselves to be as far apart as possible - All electrons in a multiple bond must lie in the same direction, hence double and triple bonds count as 1 pair of electrons. The theory refer to negative charge centres ( = pairs of electrons)

29. The 5 basic molecular shapes show the arrangement of the electron pairs (charge centres) that result in minimum repulsion between the bonding and lone pairs of electrons.

30. Order of repulsion : lone pair-lone pair > lone pair- bonding pair > bonding pair – bonding pair Methane, CH4 Bond angle is 109.50 Ammonia, NH3 Greater repulsion by lone pair of electrons. Bond angle is smaller than 109.50(1050) Water, H2O Even greater repulsion by two lone pair of electrons. Bond angle is even smaller (1050)

31. Polarity of molecules • A dipole is established when two electrical charge of opposite sign are separated by a small distance. • Polar molecules are formed between 2 atoms of different electronegativities. • Polarity of a molecule depends on the • the relative electronegativities of the atoms in the molecule and • the shape of the molecule

32. A molecule with atoms of different electronegativities may be non-polar even though there are polar bonds in the molecule as the dipoles may cancel each other and the overall dipole moment is 0.

33. When there is an uneven spread of electrons making a covalent bond, the bond is called a polar covalent bond. When there is an uneven spread of electrons over a molecule giving an unequal spread of charge over the molecule, the molecule is said to be polar. Water is an example of a polar molecule. Like bond polarity, molecular polarity is also shown using δ- and δ+.

34. Consequently, polar molecules arise when there is a net direction of charge over the molecule: polar bonds arranged in such a way as to give a net direction of charge. However, there are some instances when the polar bonds are arranged symmetrically so as to give zero net direction of charge; this is a non-polar molecule. For example, carbon dioxide and carbon tetrachloride,

35. Note that non-polar bonds can never give rise to polar molecules. Some molecules have very low polarity - so low as to be regarded as non-polar,

36. Giant covalent lattice • Usually consists of a 3-D lattice of covalently bonded atoms . • The atoms can be either same like silicon and carbon (graphite and diamond) or of 2 different elements such as silicon dioxide. • Allotropes are two (or more) crystalline forms of the same element, in which the atoms ( or molecules) are bonded differently.

37. Allotropes of carbon Diamond • Each C atom is tetrahedrally bonded to 4 other C atoms by single covalent bonds. • Very strong C-C covalent bonds have to be broken before melting occurs. It is very hard and has high very high melting point (~40000C) • All the electrons are used up in bonding (held tightly between) the atoms, and are not mobile. Hence, the electrons are localized, it does not conduct electricity.

38. Graphite • Each C atom is covalently bonded to only 3 other C atoms to give layers of hexagonal rings. Weak van der Waals’ force operates between the layers, due to the large surface area. • The layers can slide over each other so it is an excellent lubricant . • Each C has a spare electron which become delocalized along the plane. Hence, graphite is a good conductor of electricity.

39. Fullerene • 60 C atoms are arranged in hexagons and pentagons to give a geodesic spherical structure similar to a football. • Following the discover of Buckminsterfullerene , many other similar carbon molecules have been isolated. • This has led to a new branch of science called nanotechnology.

40. Intermolecular forces Van der Waals’ forces • Chance charge separation - Electrons can at any moment be unevenly spread producing a temporary instantaneous (fluctuating) dipole. • An instantaneous dipole can induce another dipole in a neighbouring particle resulting in a weak attraction between the two particles. • The forces of attraction between temporary or induced dipoles are known as Van der Waals’ forces. • Van der Waals’ forces increases with increasing mass.

42. Dipole-dipole forces • Polar molecules are attracted to each other by electrostatic forces. • Although still relatively weak the attraction is stronger than van der Waals’ forces but weaker than ionic or covalent bonds. • For polar substances with similar relative molecular masses, the higher the dipole moment, the stronger the dipole-dipole attractions and the higher the boiling points.

43. Hydrogen Bonding • If two molecules of hydrogen fluoride are close to one another, the H atom of one molecule will be attracted to the F of the other molecule, because of the electrostatic attraction between the partial charge on the hydrogen atom and the particla charge on the fluorine atom. • This charge separation or dipole exists because F is more electronegative than H. • The electrostatic attraction that holds the H atom of one molecule to the fluorine of another molecule is an example of hydrogen bond.

44. Hydrogen Bonding • The essential requirement for its formation are a H atom directly attached to O, N or F and a lone pair of electrons on the electronegative atom. • In ammonia molecule, the N atom h 1 lone pair of electrons. Each NH3 molecule van form 1 H bond. N is larger and < electronegative than F, hence the H bonding is weaker than that formed by HF

45. Hydrogen Bonding • Each the water molecule has 2 lone pairs of electrons which can form H bonds with 2 other water molecules. • The collective strength of the H bonds in water is greater than the strength of the H bonds in HF because each O atom (with 2 lone pairs) in the water molecule can form 2 H bonds with 2 other water molecules, whereas each F atom in HF molecule can only form 1 H bond with another HF molecule.

46. Effects of H bonding on physical properties Hydrogen bonding affects • the boiling points of water, ammonia, hydorgen fluoride and other molecules • the solubility of simple covalent molecules such as ammonia, methanol and ethanoic acid in water • the density of water and ice. • the viscosity of liquids, e.g. the alcohols.

47. The boiling pts of hydrides

48. Metallic bonding • The valence electrons in metals become detached from the individual atoms so that metals consist of a close packed lattice of positive ions in a sea of delocalized electrons. • A metallic bond is the attraction that two neighbouring positive ions have for the dolocalized electrons between them.

49. Metals are • malleable, that is, they can be bent and reshaped under pressure. • ductile, which means they can be drawn out into a wire. Explanation • The valence electrons do not belong to any particular atom, hence, if sufficient force is applied to the metal, 1 layer of metals can slide over another without disrupting the metallic bonding.

50. Explanation • The valence electrons do not belong to any particular atom, hence, if sufficient force is applied to the metal, 1 layer of metals can slide over another without disrupting the metallic bonding. • The metallic bonding in metal is strong and flexible and so metals can be hammered into thin sheets (malleability) or drawn into lonng wires (ductility) without breaking. If atoms of other elements are added by alloying, the layers of ions will not slide over each other so readily. The alloy is thus less malleable and ductile and consequently harder and stronger.