Slide 1

Slide 2 Slide 3 Slide 4 DiamondSilicon dioxide Graphite Slide 5 Ionic bonding typically occurs between metal and non-metal. E.g. Barium fluoride, BaF 2 The reactivity of metals and non-metals can be assessed using electro-negativity Electro-negativity ability of an atom in a covalent bond to attract ___________________________ to itself. shared paired of electrons Slide 6 Metals generally have low electronegativity values, while non-metals have relatively high electronegativity values. Fluorine, which has the greatest attraction for electrons in bond-forming situations (highest E value). Slide 7 Type of BondElectronegativity Difference Ionic2.0 4.0 Polar Covalent0.5 1.9 Nonpolar covalent0 0.4 Slide 8 Polar covalent bonds are covalent bonds with ionic character. Electrons are not shared. E.g. Na + Cl -, electron is transferred. Electrons are equally shared. E.g.Cl-Cl Ionic bond Non polar (pure covalent) bond Polar covalent bond Electrons are not equally shared. E.g. Atoms have different electronegativity values Slide 9 What type of bond is the following? (a) N (3.0) and H (2.1) (b) H (2.1) and H(2.1) (c) Ca(1.0) and Cl(3.0) (d) Al (1.5) and Cl(3.0) (e) H (2.1) and F(4.0) Type of BondElectronegativity Difference Ionic2.0 4.0 Polar Covalent0.5 1.9 Nonpolar covalent0 0.4 Slide 10 No bond is purely ionic or covalent..they have a little bit of both characters. When there is unequal sharing of electrons a dipole exists. Dipole - is a molecule that has 2 poles or regions with opposite charges. - is represented by a dipole arrow towards the more negative end. Slide 11 Intermolecular attractions are weaker than ionic, covalent, and metallic bonds Besides ionic, metallic, and covalent bonds, there are also attractions between molecules There are 2 main types of attractions between molecules: Van der Waals and Hydrogen Slide 12 Van der Waals forces consists of the two weak attractions between molecules 1. dipole interactions polar molecules attracted to one another 2. dispersion forces caused by the motion of electrons (weakest of all forces) Slide 13 Hydrogen Bonds are forces where a hydrogen atom is weakly attracted to an unshared electron pair of another atom Slide 14 This other atom may be in the same molecule or in a nearby molecule, but always has to include hydrogen. INTERMOLECULAR HYDROGEN BONDING Slide 15 Hydrogen Bonds have about 5% of the strength of an average covalent bond Hydrogen Bond is the strongest of all intermolecular forces INTRAMOLECULAR HYDROGEN BONDING Slide 16 A Network Solid contains atoms that are all covalently bonded to each other A few solids that consist of molecules do not melt until the temperature reaches 1000C or higher called network solids (Example: diamond, silicon carbide) Melting a network solid would require breaking bonds throughout the solid (which is difficult to do) http://library.thinkquest.org/C006669/data/Chem/bonding/inter.html Slide 17 The bonding pair of electrons spends most of its time between the two atomic nuclei. screening the positive charges from one another and enabling the nuclei to come closer together. Negative charge on the electron pair attracts both nuclei and holds them together in a covalent bond. When two atoms are chemically bonded, the two atoms close together have less energy and therefore are more stable than when separated. Energy is given off by the atoms to form a bond, and energy must be supplied (absorbed) to break the bond. A covalent bond is the result of electrostatic attraction between the nuclei of the 2 atoms and the pair of shared electrons. Slide 18 Slide 19 Slide 20 Number of bonds elements prefers depending on the number of valence electrons. In general - X Family # Covalent Bonds* Halogens F, Br, Cl, I Calcogens O, S Nitrogen N, P Carbon C, Si O N C 1 bond often 2 bond often 3 bond often 4 bond always Slide 21 When compounds are formed they tend to follow the Octet Rule. Octet Rule: Atoms will share electrons (e - ) until it is surrounded by eight valence electrons. Rules of the (VSEPR) game- i ) O.R. works mostly for second period elements. Many exceptions especially with 3rd period elements (d-orbitals) ii) H prefers 2 e - (electron deficient) iii) :C:N::O::F: 4 unpaired3unpaired2unpaired1unpaired up = unpaired e- 4 bonds3 bonds2 bonds1 bond O=C=ON NO = OF - F iv) H & F are terminal in the structural formula (Never central)........ Slide 22 The atomic arrangement for a molecule is usually given. CH 2 ClF HNO 3 CH 3 COOHH 2 SeH 2 SO 4 O 3 HC F Cl H H N OO O H O S O H O O OO O H C C O H H H O HSe H In general when there is a single central atom in the molecule, CH 2 ClF, SeCl 2, O 3 (CO 2, NH 3, PO 4 3- ), the central atom is the first atom in the chemical formula. Except when the first atom in the chemical formula is Hydrogen (H) or fluorine (F). In which case the central atom is the second atom in the chemical formula. Find the central atom for the following: 1) H 2 Oa) Hb) O2) PCl 3 a) Pb) Cl 3) SO 3 a) Sb) O4) CO 3 2- a) Cb) O 5) BeH 2 a) Beb) H6) IO 3- a) Ib) O Slide 23 Count the no. of valence electrons. (i) If the species has a n charge, add n to the electrons (ii) If the species has a +n charge, subtract n from the electrons Draw a skeletal structure. (i) If C is present, place C at the centre. (ii) If C is not present, place the LEAST electronegative atom at the centre. Note: H is never the in the center. Complete the octets of the outer atoms (except for H) by adding lone pairs of electrons (including the 2 electrons shared with the central atom) If there are any electrons left over, place them on the central atom as lone pairs. If the central atom does not have a complete octet, rearrange lone pairs on the outer atoms to form double bonds between the central and outer atoms. Continue doing until the central atoms octet is satisfied. If the species is charged, place it inside brackets and write the charge outside the brackets. Slide 24 Draw Lewis structures of the following molecules: (a) H 2 O, NH 3, CO 2, OCl 2, PCl 4 + (b) SO 2, NO +, OCN -, COF 2, CO 3 2-, NO 2 -, O 3 (c) BeCl 2, BH 3, PCl 5 In which of the above obey the octet rule? Bond pair _____ Lone pair.. Slide 25 1.Molecules with an odd number of electrons 2.Molecules in which an atom has less than an octet 3.Molecules in which an atom has more than an octet Slide 26 1. Odd Number of Electrons NO Number of valence electrons = 11 NO 2 Number of valence electrons = 17 O2O2 Resonance occurs when more than one valid Lewis structure can be written for a particular molecule (i.e. rearrange electrons) Molecules and atoms which are neutral (contain no formal charge) and with an unpaired electron are called Radicals Slide 27 2. Less than an Octet Includes Lewis acids such as halides of B, Al and compounds of Be BCl 3 Group 3A atom only has six electrons around it However, Lewis acids accept a pair of electrons readily from Lewis bases to establish a stable octet Slide 28 AlX 3 Aluminium chloride is an ionic solid in which Al 3+ is surrounded by six Cl -. However, it sublimes at 192C to vapour Al 2 Cl 6 molecules B2H6B2H6 A Lewis structure cannot be written for diborane. This is explained by a three-centre bond single electron is delocalized over a B-H-B Slide 29 Octet Rule Always Applies to the Second Period = n 2 ; number of orbitals ---orbitals cannot hold more than two electrons 2s, 2p x, 2p y, 2p z ---orbitals cannot hold more than two electrons Ne [He]; 2s 2, 2p x 2, 2p y 2, 2p z 2 n = 2 n = 3 Slide 30 Ar [Ne]; 3s 2, 3p x 2, 3p y 2, 3p z 2 3d 0 3d 0 3d 0 3d 0 3d 0 n = 3 Slide 31 3. More than an Octet PCl 5 Elements from the third Period and beyond, have ns, np and unfilled nd orbitals which can be used in bonding P : (Ne) 3s 2 3p 3 3d 0 Number of valence electrons = 5 + (5 x 7) = 40 10 electrons around the phosphorus SF 4 S : (Ne) 3s 2 3p 4 3d 0 Number of valence electrons = 6 + (4 x 7) = 34 The Larger the central atom, the more atoms you can bond to it usually small atoms such as F, Cl and O allow central atoms such as P and S to expand their valency. Slide 32 Bond Strength Triple bonds > Double bonds > Single bonds Page 93 Slide 33 (1)The attraction between the 2 nuclei for 3 electron pairs in a triple bond is > that for 2 electron pairs in a double bond which is > than that for 1 electron pair in a single bond. (2) Triple bonds are shorter due to greater attraction between the bonding electrons and the nuclei with more electrons in the bond. Slide 34 Strength Triple bonds > Double bonds > Single bonds Length Single bonds > Double bonds > Triple bonds Bond TypeLength (nm) Strength (kJmol -1 ) C-C0.154348 C=C0.134612 CCCC 0.120837 Slide 35 In some molecules and polyatomic ions, both electrons to be shared come from the same atom forming 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. Page 94How do you draw the Lewis structure? Slide 36 Something similar happens. A hydrogen ion (H + ) is transferred from the chlorine to one of the lone pairs on the oxygen atom. The H 3 O + ion is variously called the hydroxonium ion. Slide 37 Other examples: The reaction between ammonia and boron trifluoride, BF 3 In BF 3, there are only 6 electrons in the outer shell of boron. There is space for the B to accept a pair of electrons. Slide 38 Due to difference in electronegativity value between the 2 atoms in the bond. Unequal distribution of electron density results in small charges on the atoms ( + and - ) Example A dipole is established when two electrical charge of opposite sign are separated by a small distance. Dipole moment ElementFONClCH Electronegativity4.03.5 3.02.52.1 Slide 39 A molecule can possess polar bonds and still non-polar. Check the Geometry of the molecule: The polar bonds are arranged symmetrically so as to give zero net direction of charge. i.e. Overall dipoles cancel so that there is no overall dipole. Slide 40 Non-polar Covalent bond No difference in electronegativity value bond consists of 2 ____________ atoms. _______ net charge. Examples : Polar Covalent bond Due to the difference in electronegativity value bond consists of 2 ____________ atoms. _______ net charge. Examples: Slide 41 In the water molecule, (i) O-H bonds are significantly polar (ii) The bent structure makes the distribution of those polar bonds asymmetrical. Slide 42 Some molecules have very low polarity - so low as to be regarded as non-polar, Slide 43 Name of moleculeFormulaPolarity of molecule Hydrogen chlorideHClPolar WaterH2OH2OPolar AmmoniaNH 3 Polar BenzeneC6H6C6H6 Non-polar Boron trichlorideBCl 3 Non-polar MethaneCH 4 Non-polar BromobenzeneC 6 H 5 BrPolar Carbon dioxideCO 2 Non-polar Sulfur dioxideSO 2 Polar TetrachloromethaneCCl 4 Non-polar For CO 2 each C-O bond is polar since O is more electronegative than C. Why is the molecule non-polar? Slide 44 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 Slide 45 Find the number of electron pairs / charge centres in the valence shell of the central atom. Electron pairs / charge centres repel each other to the positions of minimum energy in order to gain maximum stability. Pairs forming a double or triple bond act as a single bond Non-bonding pairs repel more than bonding pairs. Slide 46 Slide 47 Slide 48 Methane (CH 4 ) tetrahedral Ammonia (NH 3 ) pyramidal Water (H 2 O) bent Carbon Dioxide (CO 2 ) - linear Slide 49 Slide 50 Valence Shell Electron-Pair Repulsion Theory (VSEPR) Procedure 1. Sum the total Number of Valence Electrons Drawing the Lewis Structure 2.The atom usually written first in the chemical formula is the Central atom in the Lewis structure 3. Complete the octet bonded to the Central atom. However, elements in the third row have empty d-orbitals which can be used for bonding. 4. If there are not enough electrons to give the central atom an octet try multiple bonds. Predicting the Shape of the Molecule 5. Sum the Number of Electron Domains around the Central Atom in the Lewis Structure; Single = Double = Triple Bonds = Non-Bonding Lone Pair of Electrons = One Electron Domain 6. From the Total Number of Electron Domains, Predict the Geometry and Bond Angle(s); 2 (Linear = 180); 3 (Trigonal Planar = 120); 4 (Tetrahedral = 109.5); 5 (Trigonal Bipyramidal = 120 and 90); 6 (Octahedral = 90) 7. Lone Pair Electron Domains exert a greater repulsive force than Bonding Domains. Electron Domains of Multiple Bonds exert a greater repulsive force than Single Bonds. Thus they tend to compress the bond angle. Slide 51 Lone pairs are held closer to the nucleus than the bonding pairs. The distance between the lone pair electrons and the bonding pairs of electrons is shorter than the distance between the bonding pairs to each other. Repulsion due to lone pairs causes the bond angles to become smaller Slide 52 Order of repulsion : lone pair-lone pair > lone pair- bonding pair > bonding pair bonding pair Ammonia, NH 3 Greater repulsion by lone pair of electrons. Bond angle is smaller than 109.5 0 (105 0 ) Water, H 2 O Even greater repulsion by two lone pair of electrons. Bond angle is even smaller (105 0 ) Methane, CH 4 Bond angle is 109.5 0 Slide 53 Consider NH 4 +, H 3 O +, NO 2 - NH 4 + As the 4 negative charge centres repel each other and take up positions in space to be as far apart as possible, the electron pairs are distributed in tetrahedral arrangement. H3O+H3O+ With one lone pair of electrons, the actual structure is trigonal pyramid with a bond angle of 107 0 for H-O-H bond. Read page 104 Slide 54 NO 2 - As the 3 negative charge centres repel each other and take up positions in space to be as far apart as possible, the electron pairs are distributed in trigonal planar arrangement. With one lone pair of electrons, the actual structure of the ion is bent with a bond angle of about 117 0 for O-N-O bond. Slide 55 Consider N 2 H 4, C 2 H 2 Slide 56 Forces between molecules. Does not exist in giant structure (ionic cpds, metals & giant covalent structure) http://chemtools.chem.soton.ac.uk/projects/emalaria/index.php?page=13 Slide 57 Intermolecular Forces: are generally much weaker than covalent or ionic bonds. Less energy is thus required to vaporize a liquid or melt a solid. Boiling points can be used to reflect the strengths of intermolecular forces (the higher the Bpt, the stronger the forces) Hydrogen Bonding : the attractive force between hydrogen in a polar bond (particularly H-F, H-O, H-N bond) and an unshared electron pair on a nearby small electronegative atom or ion Very polar bond in H-F. The other hydrogen halides dont form hydrogen bonds, since H-X bond is less polar. As well as that, their lone pairs are at higher energy levels. That makes the lone pairs bigger, and so they don't carry such an intensely concentrated negative charge for the hydrogens to be attracted to. Slide 58 Hydrogen Bonding & Water Slide 59 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 (London Dispersion Forces). Van der Waals forces increases with increasing mass. Slide 60 London Dispersion Forces significant only when molecules are close to each other Prof. Fritz London Due to electron repulsion, a temporary dipole on one atom can induce a similar dipole on a neighboring atom Slide 61 The ease with which an external electric field can induce a dipole (alter the electron distribution) with a molecule is referred to as the "polarizability" of that molecule The greater the polarizability of a molecule the easier it is to induce a momentary dipole and the stronger the dispersion forces Larger molecules tend to have greater polarizability Their electrons are further away from the nucleus (any asymmetric distribution produces a larger dipole due to larger charge separation) The number of electrons is greater (higher probability of asymmetric distribution) thus, dispersion forces tend to increase with increasing molecular mass Dispersion forces are also present between polar/non-polar and polar/polar molecules (i.e. between all molecules) Slide 62 The strength of the intermolecular forces determines how easily the molecules will separate and hence the melting and boiling points. Why is there an increasing boiling points of the noble gases as you go down the group? The boiling points of the noble gases are Helium-269C Neon-246C Argon-186C Krypton-152C Xenon-108C Radon-62C Slide 63 Because electrons are always moving around very quickly, the charges switch around all the time. the more electrons in a molecule / atom, the stronger these Van der Waals or London forces are. This is seen in the increasing boiling points of the noble gases as you go down the group. The boiling points of the noble gases are Helium-269C Neon-246C Argon-186C Krypton-152C Xenon-108C Radon-62C Slide 64 Cl 2 : gas I 2 : solid Iodine molecule is made up of larger atoms with more electrons compared to chlorine. With more electrons moving around, the temporary dipole will be larger. The larger atoms in the molecule means that the valence electrons are less strongly held, Hence the induced dipoles will be larger. Slide 65 van der Waals forces are present between covalent molecules with no H atom attached to N, O or F. E.g. van der Waals forces are present between HCl molecules. Theres also other intermolecular forces beween the molecules : permanent dipole - permanent dipole Slide 66 These intermolecular forces between polar molecules are stronger than between non- polar molecules. (all things being equal) For polar substances with similar RMM, the higher the dipole moment, the stronger the dipole-dipole attractions and the higher the boiling points. Slide 67 Propane (C 3 H 8 ) and ethanal (CH 3 CHO) both with RMM = 44 Ethanal has a higher bp. Ethanal - is a polar molecule - has stronger intermolecular forces (van der waals & dipole-dipole interactions) between the molecules of ethanal than between the propane molecules. Read further for a few exceptions It is not true that polar molecules have stronger intermolecular forces and hence higher bp than non-polar molecules. Non-polar molecules with higher RMM might have higher bp. Slide 68 Boiling points increase for polar molecules of similar mass, but increasing dipole: Substance Molecular Mass (amu) Dipole moment, u (D) Boiling Point (K) Propane C 3 H 8 440.1231 Dimethyl ether CH 3 OCH 3 461.3248 Methyl chloride CH 3 Cl 502.0249 Acetaldehyde CH 3 CHO 442.7294 Acetonitrile CH 3 CN 413.9355 Slide 69 A hydrogen bond is a weak type of force that forms a special type of dipole-dipole attraction which occurs when a hydrogen atom bonded to a strongly electronegative atom exists in the vicinity of another electronegative atom with a lone pair of electrons. These bonds are generally stronger than ordinary dipole-dipole and dispersion forces, but weaker than true covalent and ionic bonds. Hydrogen bonding is present between covalent molecules with H atoms attached to O, N and F Slide 70 Strength of H bond The larger the electronegativity of H and the other atom (N, O or F), the stronger the H bond. Strength F > O > N Number of them that can be formed between neighbouring molecules Slide 71 Although the strength is such F > O > N, HF can only form 1 H bond to 1 neighbour. H 2 O can form 2, thus promoting more intermolecular interactions. 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. Slide 72 Ammonia molecule has 3 N-H bonds. N is larger and < electronegative than F, has far weaker H bonds due to lower electron density on the N atom (only 1 lone pair) compared to O and F. Slide 73 When the RMM is large, we expect the boiling point to be high because larger molecules have more space for electron distribution and more possibilities for instantaneous dipole moment. However, CompoundRMMBoiling pt (K) H2OH2O18373 HF20292.5 NH 3 17239.8 H2SH2S34212 HCl36.4187 PH 3 34185.2 Greater intermolecular force because H 2 O, HF, NH 3 all exhibit hydrogen bonding. tend to have higher viscosity than those that do not have H bond. Substances which have multiple H bonds exhibit even higher Viscosity. Slide 74 Electronegativity Cannot occur without significant electronegativity difference between H and the atom it is boded to. E.g. Both PH 3 and NH 3 have trigonal pyramidal shape but only NH 3 has H bonding. Atomic size When the radii of the 2 atoms differ greatly, their nuclei cannot achieve close proximity when they interact resulting in weak interaction. Slide 75 Is there any hydrogen bonding between the molecules if CH 3 F? H H H C F H C F H H H is not joined directly to F in each molecule, hence no hydrogen bonding between the molecules. Slide 76 Is there any hydrogen bonding between the molecules of ethanol? Slide 77 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. Slide 78 Group 4A hydrides Groups 4, 5, 6A hydrides Van der Waals forces are made of dipole-dipole and London dispersion forces Slide 79 Slide 80 Recall: In ionic compounds, the ions are held together by strong ionic bonds in a giant ionic lattice. In simple covalent molecules, the attractive forces between the molecules are known as intermolecular forces or van der Waals forces, which is weaker than the ionic bonds. In giant covalent molecules, the atoms are held together by strong covalent bonds in a giant covalent lattice. Slide 81 Types of Covalent Substances Covalent substances can be divided into 2 categories as shown in the table below: Slide 82 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. Slide 83 DiamondGraphite Giant Molecular Structure Giant covalent structure Each carbon atom is covalently bonded to 4 other carbon atoms in a tetrahedral arrangement. Rings of C atoms in layers Each carbon atom is covalently bonded to 3 other carbon atoms in each layer. The fourth electron is not used in chemical bonding. This gives hexagonal rings of six atoms that join together to from flat layers that are held together by van der Waals forces. Slide 84 C60 Fullerene Molecular structure consists of individual C60 molecules with covalent bonds within the molecule. 60 C atoms are arranged in hexagons and pentagons. Van der Waals forces between molecules. Slide 85 DiamondGraphite C60 fullerene Melting/ Boiling point 1.Hard 2.Very high melting point and boiling point Strong covalent bond in the giant structure between C atoms. A lot of energy needed to break them when melts. 1.Soft and slippery 2.High melting point and boiling point Layers can slip pass each other. Covalent bond in layers must be broken when it is melted. Van der Waals force is easier to break. 1.Soft and slippery 2.Low melting point and boiling point Few covalent bonds hold the molecules together but only weak Vander Waals forces between molecules. Only Van der Waal's interactions have to be broken for melting. Properties Slide 86 DiamondGraphite C60 fullerene Electrical conductivity 1.Non-conductor of electricity No free electrons to conduct electricity. 1.Good conductor of electricity Free electron is not used in bonding, able to move within the layers and free to conduct electricity. 1.Non-conductor of electricity No movement of electrons available from one molecule to the next. Properties Slide 87 DiamondGraphite C60 fullerene Solubility 1. Insoluble in water There are only very weak Van der Waal's attractions between the carbon atoms and the water molecules whereas the carbon atoms are bonded very tightly to one another. 1.Insoluble in water There are only very weak Van der Waal's attractions between the carbon atoms and the water molecules whereas the carbon atoms are bonded very tightly to one another. 1.in water but soluble in methyl benzene There are only very weak Van der Waal's attractions between the carbon atoms and the water molecules whereas the carbon atoms are bonded very tightly to one another in the molecules. Uses1.Gemstones 2.As tips of grinding, cutting and polishing tools. 1.In pencils 2.As a dry lubricant 3.Brushes for electric motors Properties Slide 88 4Metallic Bonds Metals consist of positive ions surrounded by a 'sea of moving electrons'. The negative 'sea of electrons' attracts all the positive ions and cements everything together. Metallic bonds are the results of the strong forces of attraction between the negative electrons and the positive ions. Hence, metals have high melting points and high boiling points. Slide 89 Physical Properties of Metals Physical Properties ExplanationDiagram High DensityThe close packing of atoms explains why most metals have a high density Good Conductor of Electricity Metals are good conductors of electricity in the solid and molten state. This is due to the sea of delocalized electrons. Slide 90 Physical Properties ExplanationDiagram Malleable and Ductile Metals are malleable (can be bent or flattened) and ductile (can be drawn into wires) because the layers of metal ions can slide over each other when a force is applied. High melting and boiling point Metallic bonds are strong bonds. Except for: Mercury has a low melting point and is a liquid at room temperature. Similarly, sodium and potassium have low melting and boiling point. Slide 91 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. Slide 92 Like tends to dissolve like. Polar substances tend to dissolve in polar solvents, such as water, whereas non-polar substances tend to dissolve in non-polar solvents, such as heptane or tetrachloromethane. Organic molecules often contain a polar head and a non-polar carbon chain tail. As the non- polar carbon chain length increases in an homologous series the molecules become less soluble in water. Ethanol is a good solvent for other substances as it contains both polar and non-polar ends. Slide 93 Water will mix with polar liquids such as ethanol. The oppositely charged ends of the different molecules attract one another forming hydrogen bonds. Slide 94 Gases are generally slightly soluble in water. A small number of gases are highly soluble because they react with water to release ions. Example, SO 2 (g) + H 2 O (g) H + (aq) + HSO 3 - (aq) This solution is known as sulfurous acid, a major component of acid rain Slide 95 Slide 96 Slide 97