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CHEMICAL BONDING INTRODUCTION Why do atoms combine to form compounds? Elements that do not have stable electron configuration of noble-gas attain such a configuration through chemical reactions. During the process their outer shell either holds eight electrons (octet) or two (doublet) electrons. The process of attaining stable electron configuration is called chemical bonding. Chemical bonding Chemical bonding is the chemical process in which atoms of elements combine to form molecules or compounds. Types of Chemical Bonding There are three types of bonding, namely: (i) Ionic or Electrovalent Bonding. (ii) Covalent Bonding and. (iii) Metallic bonding. 1. Ionic (Electrovalent) Bonding Ionic bonding is a type of bonding where electrons are completely transferred from one atom to another atom or group. The strong force of attraction between the oppositely charged positive and negative ions results in an ionic bond or electrovalent bond. In an ionic structure, the ions are arranged in a regular repeating pattern called lattice. Facts about ionic bonding The electron donor loses electron(s) equal to its valence. The electron(s) move to the outer shell of the electron accepter (nonmetallic atom). Both atoms acquire the electronic structure of a noble gas. In ionic bonding, oppositely charged ions cations and anions are formed. The ions attract each other by electromagnetic force and form ionic bonds. Dot-and-cross diagrams We use dots and crosses to show the electronic configuration of the metallic and non-metallic ions so as to help us keep track of where the electrons come from and where they move to. A dot-and-cross diagram shows: (i) the outer electron shell only (ii) that the charge of the ion is spread evenly, by using square brackets (iii) the charge on each ion, written at the top right-hand corner of the square brackets. NB: When drawing a dot-and cross diagram for an ionic compound, we draw the outer electron shell of the metal ion without any electrons. Some examples of dot-and-cross diagrams (i) Sodium chloride When sodium reacts with chlorine to form sodium chloride the one electron in the outer shell of each sodium atom are transferred to the incompletely filled orbitals of a chlorine atom. By losing one electron, each sodium atom achieves the electron configuration [2,8] of neon. By gaining one electron, each chlorine atom achieves the electron configuration [2,8,8] of argon. The sodium ion (Na+) and the chlorine ion (Cl-) attract to form sodium chloride. + Na Cl Na 2,8,1 2,8,7 [2,8]+ - Cl [2,8,8]- (ii) Calcium oxide When magnesium reacts with oxygen to form magnesium oxide the two electrons in the outer shell of each magnesium atom are transferred to the incompletely filled orbitals of an oxygen atom. By losing and gaining two electrons, both magnesium atom and oxygen atom respectively achieve the electron configuration [2,8] , noble-gas configuration of neon. The magnesium ion (Mg2+) and the oxygen ion (O2-) attract each other to form magnesium oxide. 2- 2+ Mg O 2,8,2 2,6 Mg [2,8]2+ O [2,8]2+ (iii) Calcium chloride Each calcium atom has two electrons in its outer shell. The two electrons are transferred to two chlorine atoms. By losing two electrons, each calcium atom achieves the electron configuration [2,8,8] noble-gas configuration of argon. The two chlorine atoms each gains one electron and achieves the electron configuration [2,8,8], noble-gas configuration of argon. The calcium ion (Ca2+) and the two chlorine ions (Cl-) combine to form calcium chloride. - Cl 2+ Ca 2,8,7 Cl 2,8,8,2 2,8,7 Ca Cl [2,8,8]- [2,8,8]2+ Cl [2,8,8]- 2. Covalent bonding Covalent bonding is the type of bonding in which one or more pairs of electrons are shared by the two atoms. A covalent bonding occurs between non-metallic atoms. Dot-and-cross diagrams for covalent bonding When drawing the arrangement of electrons in a molecule, we (i) Use ‘dot’ for electrons from one of the atoms and ‘cross’ for the electrons from the other atom. (ii) If there are more than two types of atom we can use additional symbols such as a small circle or a small triangle. (iii) Draw the outer electrons in pairs, to emphasise the number of bond pairs and the number of lone pairs. (iv) Put the pair of shared in the region of the overlapping shell. Types of covalent bonds There are three types of covalent bonds, namely; (i) Single covalent bond. (ii) Multiple covalent bonds. - Double covalent bond - Double-double covalent bond Co-ordinate or dative covalent bond (a) Single covalent Bonds A single covalent bond or a bond pair shared pair of electrons. A single covalent bond is represented by a single line between the atoms. Examples of single covalent bonds among others are seen in; (i) Hydrogen molecule, (ii) Chlorine molecule, (iii) Water molecule (iv) Hydrogen chloride (v) Methane (i) Formation of hydrogen molecule (H2) A hydrogen molecule is formed by combination of two hydrogen atoms. The two atoms approach one another and their outer shells overlap. Each atom contributes one electron to form a single covalent bond. H (1) + H (1) H H (2) (2) H-H (ii) Formation of Chlorine molecule A chlorine molecule is formed when two chlorine atoms combine. The two chlorine atoms approach one another and their outer shells overlap. Each atom contributes one electron to form a bond pair. Cl (2,8,7) + Cl (2,8,7) Cl Cl (2,8,8) (2,8,8) Cl-Cl (iii) Formation of Hydrogen chloride molecule A hydrogen chloride molecule is formed when one hydrogen atom combines with one chlorine atom. The two atoms approach one another and their outer shells overlap. Each atom contributes one electron to form a bond pair. H (1) + Cl (2,8,7) H Cl (2) (2,8,8) H-Cl (iv) Formation of water molecule Water molecule is formed when one oxygen atom combines with two hydrogen atoms. The three atoms approach one another and their outer shells overlap. The oxygen atom contributes two electrons and each hydrogen atom contributes one electron two single covalent bonds. (v) Formation of Methane molecule Methane molecule is formed when one carbon atom combines with four hydrogen atoms. The four hydrogen atoms approach the carbon one. The outer shells overlap. Each atom contributes one electron resulting to four single covalent bonds. (2,8) 2 H (1) + O (2,6) H (2) O H (2) H 4 H (1) + C (2,4) H O H H H H C H Each hydrogen atom shares two electrons with carbon H (b) Multiple Covalent Bonds Atoms can form covalent bonds by sharing more than one pair of electrons between them. The types of multiple bonds are; (i) Double bond. E.g Oxygen (O2), Ethene (C2H4) (ii) Double-double bond E.g. Carbon dioxide (CO2). (iii) Triple bond. E.g. Nitrogen (N2) (i) Formation of oxygen molecule Oxygen molecule is formed by combination of two oxygen atoms. The two atoms approach one another and their outer shells overlap. Each atom contributes two electrons resulting to two double bonds. O (2,6) + O (2,6) O O (2,8 (2,8) O=O (ii) Formation of Ethene molecule Ethene molecule is formed by combination of two carbon atoms and four hydrogen atoms. The two carbon atoms approach one another and their outer shells overlap. Each carbon atom contributes two electrons resulting to two double bonds between the carbon atoms. Of the four hydrogen atoms, two atoms overlap with one carbon atom to form four bond pairs. H 4 H (1) + 2 C (2,4) H O H O H (iii) Formation of Carbon dioxide molecule (CO2) Carbon dioxide molecule is formed by combination of one carbon atom and two oxygen atoms. The two oxygen atoms approach the carbon atom on either side and their outer shells overlap. The carbon atom contributes all the four valence electrons and each oxygen atom contributes a pair of electrons resulting to double, double bonds. (2,6) O + O (2,6) C (2,4) O O O=C=O (iii) Formation of Nitrogen molecule Nitrogen is formed by combination of two nitrogen atoms. The two nitrogen atoms approach one another and their outer shells overlap. Each atom contributes three electrons to complete the octet stable structure resulting to a triple bond. A triple bond is represented by three short lines (≡) between the nitrogen atoms. N (2,5) + N (2,5) N N N≡N (c) Co-ordinate bonding (dative covalent bonding) A co-ordinate bond is a type of covalent bond where the pair of shared electrons is donated by one atom. For a dative covalent bonding to occur, (i) The donor atom must contain a lone pair (a pair of electrons that are not involved in bonding). (ii) The acceptor atom must have an unfilled orbital. A co-ordinate bond is represented by an arrow pointing from the donor to the recipient atom. Examples of co-ordinate covalent bonding are; - Ammonium ion. - Aluminium chloride molecule. (i) Formation of ammonium ion (NH4+) Ammonium ion is formed by combination of ammonia molecule and a hydrogen ion. The nitrogen atom in ammonia has a lone pair of electrons. It donates the lone pair of electrons to the hydrogen ion. H H H + C + + H H NH3 (g) H O H + H+ (aq) NH4+ (aq) H (ii) The formation of Aluminium chloride molecule Aluminium chloride is a covalent molecule. At high temperatures it exists as molecules with the formula AlCl3. At lower temperatures two chlorine atoms each on a separate molecule donates a lone pair of electrons to an aluminium atom on separate molecule to form a molecule Al2Cl6 . The molecule Al2Cl6 has two co-ordinate covalent bonds. Cl Cl Cl Al Cl Al + Cl Cl Cl Al Al Cl Cl Cl Cl Cl Cl Cl Al Cl Cl Al Cl Cl Shapes of Simple Molecules The shape of a refers to the three-dimensional arrangement of the atoms that constitute a molecule. It determines several properties of a substance including; reactivity, polarity, phase of matter, colour, magnetism and biological activity. The 'Electron-pair' Repulsion Theory is used to predict shapes and bond angles of simple molecules. Electron-pair Repulsion Theory The main points of electron-pair repulsion theory are: (i) (ii) In a molecule there are two types of active set of electron pairs surrounding the central atom; - the bond pair and - lone pair. The electron pairs repel one another. (iii) The repulsion forces the pairs of electrons apart until the repulsive forces are minimised. (iv) The force of repulsion between lone pairs and bond pairs is not the same. The order of repulsion The order of repulsion is as follows: lone pair-lone pair > lone pair-bond pair > bond pair-bond pair. Factors which determine the shape of a molecule The shape of a molecule depends upon total number of bond pairs and lone pairs of electrons surrounding the central atom. How to Determine a Molecular Structure 1. Draw the Lewis Structure. 2. Count the bonds and lone pairs about the central atom. NB: Count the multiple bonds as one region of electron density. 3. Determine the basic arrangement of all electron density regions (bonds + lone pairs) about the central atom. 4. Determine the molecular geometry based on resulting positions of atoms. How to predict the shapes of molecules using bond pairs only (i) If the central atom has two bond pairs and no lone pair or has two double bond, the shape is linear with bond angles of 180o. E.g. Carbon dioxide (CO2) (ii) If the central atom has three bond pairs, and no lone pair, the shape is trigonal with bond angles of 120o. E.g. Boron triflouride (BF3) (iii) If the central atom has four bond pairs and no lone pair, the shape is tetrahedral with bond angles of 109.5o E.g. Methane (CH4) (iii) If the central atom has six bond pairs and no lone pair, the shape is octahedral with bond angles of 90o E.g. Sulfur hexafluoride (SF6) How to predict the shapes of molecules using bond pairs and lone pairs (i) If the central atom has two bond pairs, and two lone pairs, the shape is bent or non-linear with bond angles of 104.5o. E.g. Water (H2O) (ii) If the central atom has three bond pairs and one lone pair, the shape is pyramidal with bond angles of 107o. E.g. Ammonia (NH3) Shapes of molecules without lone pair Examples of molecules without lone pairs are; BF3, CO2, CH4 and SF6. (i) The shape of boron triflouride (BF3) The central atom, B has three bond pairs in its outer shell. Minimizing the repulsion causes the boron triflouride molecule to have a trigonal planar shape with a bond angle of 120o. F B F F (ii) The shape of Carbon dioxide (CO2) The central atom, C atom has four bond pairs in its outer shell. The molecule has two double bonds. Minimizing the repulsion causes the carbon dioxide molecule to have a linear shape with a bond angle is 180o. O O (iii) The shape of Methane (CH4) The central atom, C atom has four bond pairs in its outer shell. Minimizing the repulsion causes the bond pairs arrange themselves about the C atom resulting to a tetrahedral shape with a bond angle of 109.5o. H H O H H H H C H H (iv) The shape of Sulfur hexaflouride (SF6) The central atom, S atom has six bond pairs in its outer shell. Minimizing the repulsion causes the bond pairs to arrange themselves about the S atom resulting to a octahedral shape with a bond angle of 90o. The 6 F atoms are located at the corners of the octahedron. F F F S F F F Shapes of molecules with lone pair Examples of molecules without lone pairs are; H2O and NH3. (i) The shape of water (H2O) The central atom, O has two bond pairs and two lone pairs in its outer shell. Minimizing the repulsion causes the water molecule to have a bent or non-linear shape with a bond angle is 104.5o. (ii) The shape of Ammonia (NH3) The central atom, N atom has three bond pairs and a lone pair in its outer shell. The lone pair repels the bond pairs and then arrange themselves about the N atom, thus resulting to a trigonal pyramid or pyramidal shape with a bond angle of 107o. H N H H Summary of shapes of molecules Molecule Bonding pair Lone pairs Bond angle Shape CO2 2 0 180o Linear BF3 3 0 120o Trigonal CH4 4 0 109.5o NH3 3 1 107o H 2O 2 2 104.5o 0 90o SF6 6 Tetrahedral Pyramidal Bent Octahedral bonds and ) bonds (i) (ii) bond (Sigma bond) A sigma bond is a bond formed a hybridised p orbital (modified p orbital) overlaps linearly (end-on) with an s orbital or another modified p orbital. bond A bond is a bond formed when p orbitals overlap sideways. A single bond is drawn as two electron clouds one arising from each lobe of the p orbitals. Shapes of some organic molecules Ethane (CH4) and Ethene (C2H4) Shapes of molecules are explained in terms of the patterns of electron density found in (i) bonds and bonds. Shape of Ethane (CH4) Ethane has a displayed formula as; All the bonds in ethane are bonds. All the areas of electron density repel each other equally. This makes the H – C – H bond angles all the same (109.5o. (ii) Shape of Ethene The displayed formula of ethene is Each carbon atom in ethene uses three of its four outer electrons to form bonds. Two bonds are formed with the other carbon atom. The fourth electron from each carbon atom occupies a p orbital, which overlaps sideways with a similar p orbital on the other carbon atom to form a bond. (iii) The shape of benzene The molecular formula of benzene is C6H6. The shape of benzene is a planar regular hexagon, with bond angles of 120°. The symbol for benzene All the bonds are identical. This means that there aren't alternating double and single bonds. The electrons are delocalised. Therefore, the symbol of benzene is; NB: The hexagon shows the ring of 6 carbon atoms, each of which has one hydrogen attached. The circle represents the delocalised electrons. Formation of Benzene structure In benzene, each of the 6 carbon atoms undergoes sp2 hybridisation. Each atom uses three hybrid orbitals to form three sigma bonds, two with two carbon atoms and one with a hydrogen atom. The p orbitals in the same plane in each carbon atom overlap sideways to form 3 bonds . The electrons in the delocalised. bonds are H H H H H H H H H H H H Bond energies, Bond lengths and Bond polarities (i) Bond length Bond length or bond distance is the average distance between nuclei of two bonded atoms in a molecule. In general double bonds are shorter that single bonds. Reason: Double bonds have a greater quantity of negative charge between the two atomic nuclei. The greater force of attraction between the electrons and the nuclei pulls the atoms closer together. (ii) Bond energy Bond energy is the energy needed to break one of a given bond in a gaseous molecule. It is usually expressed in kJ mol-1. Table showing some values of bond lengths and bond energies Bond Bond Energy/(kJmol-1) Bond length/nm C-C 350 0.154 C=C 610 0.134 C-O 360 0.143 C=O 740 0.116 Relationship between Bond length and Bond energy For non-metallic atoms in the same group bonded to the same element, the bond energy decreases down the group while the bond lengths increase down the group. Example The table below shows the bond lengths and bond energies of some. Hydrogen halide Bond length/nm Bond Energy/ (kJmol-1) H - Cl 0.127 431 H - Br 0.141 366 H-I 0.161 299 (a) What is the relationship between the bond lengths and the bond energy for these hydrogen halides? (b) Suggest why the bond energy values decrease in the order HCl > HBr > HI. (c) Suggest a value for the bond length in hydrogen fluoride, HF. Electronegativity and Bond polarity (i) Electronegativity Electronegativity is the ability of a particular atom, which is covalently bonded to another atom, to attract the bond pair of electrons towards itself. The greater the value of the electronetavity, the greater is the power of an atom to attract the electrons in a covalent bond towards itself. Solution (a) The bond energies of the halides decrease down the group since the bond lengths of the halides increase down the group. (b) The atomic size of the halogens increase down the group. Therefore, the reactivity of the halogens decrease down the group. (c) 0.122 nm The pattern of Electronegativity in the Periodic Table Electronegativity; (i) General increases across a period from Group I to Group VII. Reason: Due to increasing nuclear charge and decreasing atomic radius. (ii) Decreases down or increases up each group. Reason: Due to increasing atomic radius and shielding effect. For the most electronegative elements, the order of electronegativity is; Increasing electronegativity F > O > N > Cl > Br > C > I > H Electronegativity values for some elements Element Electronegativity F 4.0 Cl 3.0 Br 2.8 S 2.5 C 2.5 H 2.1 (ii) Bond polarity The polarity of a covalent bond depends on the difference of the electronegativity values of the atoms forming the bond. The difference can make the covalent bond to be either non-polar or polar. How to predict the type of bond using electronegativity differences The difference in electronegativities between the two elements is determined. (i) If the difference is ionic. 1.7, the bond is (ii) If the difference is < 1.7 but > 0.5, the bond is polar. (iii) If the difference is < 0.5, the bond is non-polar. Non-polar Bonds When the electronegativity values of the two atoms forming a covalent bond are the same, the pair of electrons is equally shared. The electron distribution is said to be symmetric and the bond is said to be non-polar. For example, the molecules such as; H2, Cl2, F2, Br2 and O2 are non-polar molecules. Polar (dipole) bonds When a covalent bond is formed between two atoms having different electronegativity values, the more electronegative atom attracts the pair of electrons in the bond towards itself. The bond becomes polar or dipole. Facts about polar bonds (i) The centre of the positive charge does not coincide with the centre of the negative charge. (ii) The electron distribution is asymmetric. (iii) The two atoms are partially charged. - The less electronegative atom is shown with the partial charge (delta positive). - + The more electronegative atom is shown with the partial (delta negative). (iv) The degree of polarity of a molecule is measured as a dipole moment. The direction of the dipole is shown by the sign - The arrow points to the partially negatively charged end of the dipole. For example, hydrogen chloride (HCl) molecule is a polar molecule and is shown as; + H Cl 3. Metallic Bonding A metal consists of a regular array of closely packed atoms. The atoms lose their outer shell electrons into the spaces between the atoms and become positive ions. The delocalised electrons distribute themselves throughout the entire piece of the metal resulting to electrostatic attraction between the positive ions and the free moving electrons. The electrostatic force of attraction constitute metallic bond. Metallic bond is strong. Reason: The ions are held together by the strong electrostatic attraction between the positive ions and the negative charges of the delocalised electrons. Factors which determine the strength of metallic bonding The strength metallic bonding increases with; (i) Increasing positive charge on the metal ions in the lattice. (ii) Decreasing size of the metal ions in the lattice. (iii) Increasing number of mobile electrons per atom. Bond polarity in molecules containing more than two atoms We take into account; (i) The polarity of each bond. (ii) The arrangement of the bonds in the molecule. Consider the trichloromethane (CHCl3) & tetrachloromethane (CCl4) molecules. Each has four bond pairs surrounding the central carbon atom. The trichloromethane (CHCl3) molecule is a polar while the tetrachloromethane (CCl4) is non-polar. Why trichloromethane (CHCl3) molecule is polar The three C-Cl bonds are dipoles due to the difference in electronegativity. The three C-Cl dipoles point in a similar direction and the combined effect of the three dipoles is not cancelled out by the polarity of C-H bond. Since the C-H bond is virtually non-polar. The electron distribution in the molecule is asymmetric. Therefore, the molecule is polar, with the negative end towards the chlorine atoms. Why tetrachloromethane (CCl4) molecule is non-polar Tetrachloromethane has four polar C-Cl bonds pointing towards the four corners of a tetrahedron. The dipoles in each bond cancel each other. So, tetrachloromethane is non-polar. (a) H C Cl Cl (b) Cl + Cl - Trichloromethane, a polar molecule C Cl Cl Cl Tetrachloromethane, a non-polar molecule Polarity and Chemical reactivity Bond polarity influences chemical reactivity. For example, both nitrogen (N≡N) and carbon monoxide (C≡O) have triple bonds requiring a similar amount of energy to break them. But carbon monoxide is more reactive than nitrogen molecule. Explanation: Nitrogen molecule Nitrogen molecule id non-polar molecule and is fairly unreactive. Carbon monoxide molecule Carbon monoxide molecule is a polar molecule. This explains its reactivity with oxygen and its use as a reducing agent. How bond polarity helps in starting a chemical reaction Many chemical reactions are started by a reagent attacking one of the electrically charged ends of a polar molecules. For example, Chloroethane (C2H5Cl) is far more reactive than ethane (C2H6). Explanation: The Chloroethane molecule is polar. H H H C + Cl C H H Therefore, reagents such as OH- ions can attack the delta positive carbon atom of the polarised C-Cl bond. Such attack is not possible with ethane because the C-H bond is virtually non-polar. Intermolecular forces Intermolecular forces are forces between molecules. Types of intermolecular forces There are three types of intermolecular forces, namely; (i) van der Waals’ forces (ii) Permanent dipole-dipole forces (iii) Hydrogen bonding The intermolecular forces are weaker than forces within molecules due to covalent and ionic bonding. A table showing the relative strength of intermolecular forces and other bonds Bond type Bond strength/ kJmol-1 Ionic bonding in NaCl 760 O-H covalent in H2O 464 Hydrogen bonding 20 - 50 Permanent dipole-dipole force 5 - 20 Van der Waals’ forces 1 - 20 (i) Van der Waals’ forces van der Waals’ forces are very weak forces between atoms and molecules. Facts about van der Waals’ forces (i) They are due to dipoles set up between molecules. The dipoles are as a result of the electron charge clouds in non-polar molecule having more of the charge cloud on one side than the other. Such that one dipole induces a dipole on the neigbhouring molecules. (ii) Van der Waals’ forces are also called temporary dipole-induced dipole forces and dispersion forces . (iii) Van der Waals’ forces increase with; - Increasing number of electrons and protons in the molecule. - Increasing number of contact points between molecules. NB: Contact points are places where the molecules come close together. They affect the boiling points of liquids. Effect of increasing number of electrons on Enthalpy of vaporisation and boiling points Enthalpy of Vaporization is the quantity of heat that must be absorbed if a certain quantity of liquid is vaporized at a constant temperature. Both enthalpy of vaporisation and boiling points of molecules increase with increase in the number of electrons as shown in graphs below. Enthalpy of vaporisation of noble gases against number of electrons Enthalpy of vaporisation/ kJmol-1 20 Xe 10 Ar Kr Ne He 0 0 20 40 60 Number of electrons Boiling points of noble gases against number of electrons Boiling/oC -100 Xe Kr Ar -200 Ne -270 He 0 20 40 60 Number of electrons Effect of increasing number of contact on boiling points Consider the two isomers pentane, CH3CH2CH2CH2CH3 and 2,2methylpropane, (CH3)4C. The two molecules have the same number of electrons but different boiling points (b.pt of pentane = 36oC and that of 2,2-methylpropane = 10oC. Explanation: The molecules in pentane can link up besides each other so there are a large number of contact points. The van der Waals’ forces are higher, so the boiling point of pentane is higher. The molecules in 2,2-dimethylpropane are more compact and have smaller number of contact points. The van der Waals’ forces are relatively lower, so the boiling point is lower. This can be illustrated as shown below. CH2 CH2 CH3 CH2 CH2 CH3 CH2 CH3 CH2 Contact point CH3 Pentane, Boiling point 36oC CH3 CH3 C CH3 CH3 Contact point CH3 Contact point CH3 CH3 C CH3 2,2-dimethylpropane, Boiling point 10oC (ii) Permanent dipole-dipole forces Permanent dipole-dipole forces are forces of attraction between two molecules having permanent dipoles. Facts about dipole-dipole forces The molecules always have negatively and positively charged ends. The attractive force between the + charge on one molecule and the - on a neigbhouring molecule causes a weak attractive force between the molecules. For small molecules with the same number of electrons, permanent dipole-dipole forces are often stronger than van der Waals’ forces. Effect of permanent dipole-dipole forces on the boiling points of compounds Consider the boiling points of propanone (CH3COCH3, Mr = 58 and butane (CH3CH2CH2CH3, Mr = 58). The two molecules have the same relative molecular mass but different boiling points (b.pt of propanone = 56oC and that of butane = 0oC. CH2 CH3 CH2 CH3 Butane, boiling point 0oC CH3 + C=O - CH3 Propanone, boiling point 56oC Explanation: The permanent dipole-dipole forces between propanone molecules are strong enough to make this substance a liquid at room temperature and has a higher boiling point. There are only van der Waals’ forces between butane molecules. These forces are comparatively weak, so butane is a gas at room temperature and has a lower boiling point. (iii) Hydrogen bonding A hydrogen bonding is the attractive force between a hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different molecule. Conditions for hydrogen bonding to occur For hydrogen bonding to occur between two molecules we need; (i) One molecule having hydrogen atom covalently bonded to F, O or N (the three most electronegative atoms). (ii) A second molecule having F, O or N atom with a lone pair of electrons. Therefore, the average number of hydrogen bonds formed per molecule depends on; (i) The number hydrogen atoms attached to F, O or N in the molecule. (ii) The number of lone pairs present on the F, O or N atom in the molecule. Examples of compounds in which there is hydrogen bonding are; Water and Ammonia (i) Hydrogen bonding in water Water has two hydrogen atoms and two lone pairs per molecule. So water is extensively hydrogen bonded with other water molecules. It has an average of two hydrogen bonds per molecules as shown in the diagram below. + H H O H + H - H + O O - - H NB: The hydrogen bonding repeats itself continuously. (ii) Hydrogen bonding in Ammonia Ammonia has three hydrogen atoms and one lone pair per molecule. So ammonia is less extensively hydrogen bonded than water. It can form on average, only one hydrogen bond per molecules as shown in the diagram below. H H H - N + H H H N CHECK UP Draw diagrams to show hydrogen bonding between the following molecules; (a) Ethanol, C2H5OH. (b) Two hydrogen fluoride molecules. Effect of hydrogen bonding on boiling point of compounds Hydrogen bonding makes some compounds to have higher boiling points than expected. For example the graph of boiling points of halides, HF, HCl, HBr and HI, plotted against the position of the halogens in the Periodic Table, shows an irregular pattern. The boiling points of halides against their position in the Periodic Table Boiling point/oC -50 0 -50 -100 HF HCl HBr HI Explaining the shape of the graph The high boiling point of fluorine is due to the stronger intermolecular forces of hydrogen bonding between the HF molecules as a result of high electronegativity of fluorine. The rise in boiling point from HCl to HI is due to the increasing number of electrons in the halogen atoms down the group. This leads to increased van der Waals’ forces as the molecules get bigger. The graph of boiling points against relative molecular mass of other hydrides CHEP UP The table below shows the boiling points of some Group V hydrides. Hydride (a) (b) Boiling point/oC Ammonia, NH3 -33 Phosphine, PH3 -88 Arsine, AsH3 -55 Stibine, SbH3 -17 Explain the trend in the boiling points from phosphine to stibine. Explain why the boiling point of ammonia does not follow this trend. 1. Effect of hydrogen bonding on water The intensive Hydrogen bonding in water causes water to have peculiar properties. High enthalpy change of vaporisation and boiling point. The enthalpy change of vaporisation and the boiling point of water are much higher than predicted by the trend in boiling points for the other Group VI hydrides. This is because water is intensively hydrogen bonded. Enthalpy of vaporisation/kJmol-1 The boiling points of halides against their position in the Periodic Table 40 20 H 2O H2Te H 2S H2Se 10 0 Relative molecular mass (not to scale) 2. Surface tension and Viscosity (i) Surface tension Water has high surface tension compared to other liquids. Reason: The hydrogen bonds in water exert a significant downward force at the surface of the liquid causing the surface tension of water to be high. (ii) Viscosity Water has high viscosity compared to other liquids. Reason: The hydrogen bonding reduces the ability of water molecules to slide over each other, so the viscosity of water is high. 3. Ice is less dense than water Reason: There is a three-dimensional hydrogen bonded network of water molecules. This produces a rigid lattice in which each oxygen atom is surrounded by a tetrahedron of hydrogen atoms. This ‘more open’ arrangement allows the water molecules to be slightly further apart than in the liquid state. So the density of ice is less than the density of water. Bonding and Physical Properties The type of bonding between atoms, ions or molecules influences the physical properties of a substance. (a) Physical state at room temperature and pressure, Boiling point, melting point and enthalpy change of vaporisation. (i) Ionic compounds Are solid at room temperature and pressure. Reason: There are strong electrostatic forces (ionic bonds) holding the positive and negative ions together. The ions are regularly arranged in a lattice, with the oppositely charged ions close to each other. Ionic compounds have high melting point, high boiling point and high enthalpy change of vaporisation. Reason: It takes a lot of energy to overcome the strong electrostatic attractive forces. (ii) Metals Metals, apart from mercury, are solids. Most metals have high melting points and boiling points and high enthalpy changes of vaporisation. Reason: It takes a lot of energy to overcome the strong attractive forces between the positive ions and the ‘sea’ of delocalised electrons. (iii) Covalent compounds Covalently bonded substances with simple molecular structure, for example water and ammonia, are usually liquids or gases. Some are solids at room temperature and pressure. E.g. Iodine, paraffin wax, polyethene etc. They have low melting point, boiling point and low enthalpy changes of vaporisation compared with ionic compounds. Reason: The forces between the molecules are weak. (b) Solubility (i) Ionic compounds Most ionic compounds are soluble in water. Reason: Water molecules are polar and they are attracted to the ions on the surface of the ionic solid. These attractions, ion-dipole, replace the electrostatic forces between the ions and the ions go into solution. (ii) Metals Metals do not dissolve in water. However, some metals react with water. E.g. Sodium, potassium & calcium. (iii) Covalent compounds Covalently bonded substances with a simple molecular structure fall into two groups. Those that are insoluble in water. Most covalently bonded molecule are non-polar. Water molecules are not attracted to them so they are insoluble. Those that are soluble in water. Small molecules that can form hydrogen bonds with water are generally soluble. E.g. Ethanol, C2H5OH. Some covalently bonded substances react with water rather than dissolving in it. E.g. Hydrogen chloride (HCl) reacts with water to form hydrogen ions and chloride ions. (c) Electrical conductivity (i) Ionic compounds Ionic compounds do not conduct electricity when in the solid state. Reason: The ions are held strongly by the electrostatic forces of attraction and are not free to move. They contact electricity in molten and aqueous. Reason: The ions are free to move. (ii) Metals Metals conduct electricity both when solid and when molten. Reason: There are free to moving electrons. (iii) Covalent compounds Covalently bonded substances with a simple molecular structure do not conduct electricity. Reason: They have neither ions nor electrons which are free to move. END OF CHAPTER