Bonding. Metallic Ionic Covalent. (Elements) (Compounds) (Elements and Compounds)

Transcription

1 CfE Higher Chemistry Unit One Chemical Changes and Structure Chapter Four Bonding in Compounds Types Of Bonding In Compounds Bonding Metallic Ionic Covalent (Elements) (Compounds) (Elements and Compounds) Covalent Molecular Covalent Network Intramolecular Intermolecular Pure Covalent Polar Covalent (Permanent Dipoles) Dipole-Dipole London Dispersal Forces Hydrogen Bonding (Special Version of Dipole-Dipole) 1

2 Types of Bonding in Compounds Ionic Bonding Ionic bonding is an electrostatic force of attraction between positive ions of metals (and some notable non-metal complex ions) with the negative ions of non-metals ( or groups of non-metals known as complex ions, for example SO 4 2- ). For example Sodium Chloride Different elements have different degrees of attraction for bonding electrons, i.e. their electronegativity. The difference in the electronegativities between the metals and non-metal atoms results in a transfer of electrons from the metal atom (low electronegativity) to the non-metal atom(s) (high electronegativity),so creating a positive metal ion and a negative non-metal ion. Ionic bonding occurs as a result of the electrostatic force of attraction between the positive metals ions and the negative non-metal ion. Greater the difference in electronegativities, between the metal atoms and the nonmetal atoms, the greater the degree of ionic bonding. 2

3 Covalent Bonding Covalent bonding occurs in general between non-metal atoms (there are some notable exceptions, for example, titanium chloride). There are two types of covalent bonds; Pure covalent Polar covalent. Pure Covalent Bonds These bonds result from the non-metal elements in the molecule having the same electronegativities. All molecules containing only pure covalent bonds have properties associated with London Dispersal Forces. All molecular elements and a number of molecular compounds contain pure covalent bonds and are therefore gases, liquids or low melting solids at room temperature. Polar Covalent Bonds These bonds result from the non-metal atoms having different electronegativities. As a result of these polar bonds the molecule itself is polar (with the exception of symmetrical molecules). Important to note; not all molecules containing polar bonds are polar molecules a group of molecules that have polar bonds but are symmetrical in shape and therefore non polar molecules as their polar bonds cancel each other out. 3

4 Compounds with a greater degree of differences in their electronegativities are considered as most ionic, while compounds with the least degree of difference in their electronegativities are considered most covalent. In general differences in electronegativities indicate whether a compound is ionic, polar covalent or indeed pure covalent but it is not precise but the Bonding Continuum can be used to help understand these subtle differences in the types of bonding associated with different compounds. To determine the true nature of the bonding present we must consider the compounds properties. 4

5 Intermolecular forces of Attraction There are three intermolecular forces of attraction (two resulting from polar bonds). 1. London Dispersal Forces The electrons of an atom that surround the nucleus are not stationary but are constantly moving around the nucleus. This movement of electrons creates an uneven distribution of electrons (negatively charged) around the (positively charged) nucleus that results in the formation of a temporary dipole on the atom. London Dispersal Forces Dipole dipole attraction The Greek letter delta (δ) is used to denote a slight amount. Thus delta negative (δ - ) means slightly negative and delta positive (δ + ) means slightly positive. The London Dispersal Forces result from the temporary attraction between the positive end of one dipole for the negative end of another induced dipole. This is a temporary dipole-dipole electrostatic force of attraction. 5

6 2. Permanent Dipole-Dipole Forces of Attraction These forces of attraction result from the non-metal atoms in the molecule having different electronegativities. For example Hydrogen sulfide H 2 S 2.5S δ- 2.2H δ+ H δ+ ιιιιιιιιιιιιιι S δ- H δ+ H δ+ Permanent dipole Permanent Dipole-Dipole force of attraction (Much stronger than London Dispersal Forces of Attraction) 6

7 3. Hydrogen Bonding (A special type of Permanent Dipole-Dipole Forces of Attraction) This is a special form of Dipole-Dipole force of attraction in which the non-metal atom bonded to the hydrogen atom has a much greater electronegativity than hydrogen. For example water H 2 O Permanent dipole O δ- O δ- ιιιιιιιιιιι δ+ H H δ+ H δ+ H δ+ Hydrogen Bonding (Permanent Dipole-Dipole force of attraction) (Stronger than normal Dipole-Dipole Forces of Attraction) Note Hydrogen Bonding occurs when hydrogen is bonded to; Oxygen Nitrogen Flourine O H N H F H 7

8 Covalent Molecular Structures Most covalent compounds are molecular. For Example Chlorine Hydrogen Flourine Cl Cl H H F F Mpt -101 C Mpt -259 C Mpt -220 C Bpt -35 C Bpt -253 C Bpt -188 C These molecular elements are gases at room temperature as they only have weak intermolecular forces of attraction known as London Dispersal Forces between their molecules. London Dispersal Forces F F ιιιιιιιιιι F - F Pure Covalent Bonds ( Note; Pure covalent bonds as both atoms in the molecule have identical electronegativities). During melting and evaporation of these elements only the very weak intermolecular London Dispersal Bonds are broken, hence the low melting and boiling point. The intramolecular pure covalent bonds remain intact. 8

9 Polar Covalent Bonds In polar covalent bonds the bonding electrons are not shared equally as the bonding non-metal atoms have different electronegativities. The atom with the higher electronegativity has a greater share of the bonded electrons, so has a slight negative charge, (δ - ). The other bonded atom with the lower electronegativity has a slight positive charge, (δ + ). This results in a polar (dipole) bond. Examples Hydrogen chloride (H-Cl) δ + δ - H ---- Cl This is a polar molecule Polar Covalent Bond Water (H 2 O) 3.5 O δ - δ + H H δ + This is a polar molecule These polar covalent bonds are also known as permanent dipoles. 9

10 Polar Molecules A polar covalent molecule is one that has permanent charged ends (δ + and δ - ) called permanent dipoles. These dipoles enable polar molecules to bond to one another by bonds known as dipole dipole forces of attraction. Consider Hydrogen Iodide, HI This intermolecular force of attraction is known as a dipole-dipole force of attraction (bond). δ + δ - δ + δ - H ---- I ιιιιιιιιιι H ---- I Polar Covalent Bonds 10

11 Symmetrical Polar Molecules Some covalent molecules contain polar bonds but are not overall polar as they are symmetrical. For example Caron dioxide Carbontetrachloride δ - δ - Cl O C O δ + C δ + Polar covalent bonds δ - Cl Cl δ - Cl δ - As a result of their symmetrical shape, their polarities cancel one another out resulting in the molecules themselves being non-polar. Their physical properties result from London Dispersal (intermolecular) forces of attraction between their molecules, (hence both these molecules are gases at room temperature). 11

12 Note The tetrahedral shape is symmetrical Methane CH 4 Subject to London Dispersal forces H C δ - Polar bonds Symmetrical molecule H H δ + Therefore a non-polar molecule H Contains polar bonds, but as the central atom carbon is surrounded bonded to four identical hydrogen atoms with identical electronegativities then this molecule is nonpolar. Methane molecules are subject to weak London Dispersal forces and are therefore a gas at room temperature. Simple alkanes like pentane, C 5 H 12 that have polar bonds (δ C molecules as they are symmetrical. H δ + ) are non-polar H H H H H H C C C C C H H H H H H The only force of attraction associated with the pentane molecules are the intermolecular London Dispersal forces of attraction. Due to the molecular mass of pentane it is a liquid at room temperature. 12

13 Non-symmetrical polar molecules Dipole dipole forces of attraction For example Ammonia NH 3 N δ - Polar bonds (permanent dipole) Non-symmetrical molecule H H Therefore a polar molecule H δ + N δ - Dipole dipole force of attraction H H δ + H Chloroform CCl 3 H Subject to dipole-dipole attraction H C δ + Polar bonds Non-symmetrical molecule Cl Cl δ - Therefore a polar molecule Cl Boiling point 61 C, gfm 119.5g This difference between methane (a gas) and chloroform (a liquid) is an indication of the greater strength of dipole-dipole attractions compared with London Dispersal Forces of attraction. Polar molecules like chloroform CCl 3 H have polar bonds (i.e. permanent dipole) within their structure and are themselves non-symmetrical in shape. 13

14 Non-symmetrical polar molecules Hydrogen bonding Water, H 2 O Boling point 100 C Gfm 18g Liquid at room temperature Dipole-dipole intermolecular force - of attraction known as Hydrogen Bonding H - O H Polar Bonds Non-symmetrical molecule Polar molecule + Permanent dipole H = + - O H This difference between chloroform (Bpt 61 C) and water (Bpt 100 C, gfm 18g) is an indication of the much greater strength of dipole-dipole attraction that is hydrogen bonding compared with normal dipole dipole attractions. Hydrogen bonding is a special form of permanent dipole to permanent dipole interaction when these molecules contain a hydrogen atom bonded to a highly electronegative element such as fluorine, oxygen and nitrogen. 2.2 δ + H F δ δ + H O δ δ + H N δ

15 Consider Ethanol C 2 H 5 OH Gfm = 46g Bpt = 79 C Ether C 2 H 6 O gfm = 46g Bpt = -23 C H H H H H C C O H H C O C - H H H H H Polar bonds and non-symmetrical Therefore a polar molecule Dipole-dipole attraction Polar bonds and symmetrical Therefore a non-polar molecule (London Dispersal Forces of Attraction) (Hydrogen bonding) When hydrogen bonding is present, then the molecule has much higher boiling points compared with other molecules with similar molecular masses and no hydrogen bonding All alcohols (-OH functional group) are subject to hydrogen bonding. Electrical Fields Polar molecules like ethanol and water bend in an electric field whereas non-polar molecules like pentane do not. 15

16 Anomalous Physical Properties of some Hydrides Group four Hydrides CH 4, SiH 4, GeH 4 and SnH 4 These show the expected increase in boiling points with molecular size, due to increased London Dispersal forces of attraction. Groups Five, Six and Seven The boiling points for ammonia (NH 3 group 5), Water (H 2 O group 6) and Hydrogen fluoride (HF group 7) all have boiling points greater than expected. The elevated boiling points indicate a stronger intermolecular forces of attraction than expected from London Dispersal Forces or permanent dipole Permanent Dipole forces of attraction. This stronger intermolecular force of attraction is Hydrogen Bonding. 16

17 Viscosity Viscosity normally increases with increasing molecular size but molecules with hydrogen bonding show higher viscosity than expected. Miscibility Miscible liquids mix thoroughly without any visible boundary eg. water and ethanol. Immiscible liquids have a boundary between them eg. water and hexane. Hydrogen bonding helps miscibility eg. both water and ethanol contain hydrogen bonding. Other polar liquids are also often miscible with water eg. propanone and water are miscible. 17

18 Solubility Ionic lattices and polar covalent compounds tend to be soluble in water and other polar solvents due to the interaction of opposite charges. When ionic compounds dissolve in water their lattice become surrounded by polar water molecules. The negative ions are attracted to the positive ends of the water molecules and the positive ions are attracted to the negative ends of the water molecule Ions surrounded by a layer of water molecules, held by electrostatic forces of attraction are said to be hydrated. Non-polar molecules will tend to be soluble in non-polar solvents like hexane or carbon tetrachloride and insoluble in water and other polar solvents as they have no charged ends to be electrostatically attracted to the polar solvent molecules. 18

19 The Structure of Ice Normally solids are denser than the own liquids, but, ice floats on water. The intermolecular bonding associated with small covalent molecules is usually London Dispersal Forces of Attraction. In ice, the intermolecular force of attraction is hydrogen bonding. This result in a crystal lattice of water molecules that are held together by an network of hydrogen bonds. If you examine the diagram closely you will see that each water molecule is surrounded by four hydrogen bonds This arrangement not only makes the structure strong but it also spaces out the water molecules and so prevents them from packing closely together. The structure of ice results in the water molecules being less densely packed together than those of water and therefore the ice floats 19

20 Covalent Network structures These structures have an infinite three-dimensional network structure of non-metal atoms bonded together by covalent bonds. These elements have extremely high melting and boiling points. Two examples are; Silicon Dioxide (sand) The resulting structure is almost as hard as diamond. Silicon Carbide Silicon carbide has a similar network structure. It is the hardest known substance to man. It is used on abrasive wheels for cutting rocks or on grinders for sharpening metals 20