Chemistry AS Textbook

NES/Chemistry/AS Intermolecular Forces of attraction, or repulsion, act between molecules and other types of neighbouring particles, such as other molecules, atoms, or ions. - There are four types of intermolecular forces: hydrogen bonding, van der Waals' (permanent dipole interactions and instantaneous dipole interactions) and ion-dipole interactions. Hydrogen Bonds are an electrostatic force of attraction between a hydrogen atom covalently bonded to an atom of either fluorine, oxygen or nitrogen in one molecule and a fluorine, oxygen or nitrogen atom of another molecule. Electronegativity is a measure of the ability of an atom in a molecule to attract the electrons in a covalent bond. Bond Length is the distance between the centres of adjacent atoms. -In general, the longer the bond is, the weaker the bond is. Bond Polarity occurs when two bonded atoms unequally share electrons, due to differences in electronegativity. The unequal sharing of electrons within a bond leads to a one end of the bond becoming more positive and the other end becoming more negative. Dipoles occur when a positive charge and a negative charge are separated by some fixed distance. -Such as a molecule with a positive charge at one end and a negative charge at the other. Dipole Moment is the measure of the net polarity of a molecule. -Molecules with a greater difference in electronegativities will have a greater dipole moment (unless they are symmetrical). Permanent Dipoles occur when two atoms in a molecule have substantially different electronegativity: One atom attracts the bond pair electrons more than another, becoming more negative, while the other atom becomes more positive. A molecule with a permanent dipole moment is called a polar molecule. Temporary Dipoles spontaneously form in a non-polar molecule due to chance when electrons happen to be more concentrated in one place than another in a molecule, creating a instantaneous dipole. -These dipoles are temporary and weaker in effect than permanent dipoles. 50

NES/Chemistry/AS Induced Dipoles occur when a polar molecule induces a dipole in an atom, or in a non-polar molecule, by disturbing the arrangement of electrons in the non-polar species. -This is when a non-polar molecule, or atom, temporarily gain polarity from a close-by polar molecule, or ion. The polarity is lost when the polar molecule, or ion, moves away. Non-polar Molecules are molecules in which the electrons are shared equally between the nuclei. As a result, the distribution of charge is even and the force of attraction between different molecules is small. Non-polar molecules show little reactivity. -This happens either when the molecule has no bond polarity, or has bond polarity, but the molecule is symmetrical. Polar Molecules have a partial positive charge in one part of the molecule and complementary negative charge in another part. -This arises from a difference in electronegativities in the atoms in the molecule. Polarising power is the ability of a cation to distort the electron cloud around an anion. Polarisability is a measure of the degree to which an atom, or anion has had its electron cloud distorted. Dimers are molecules or molecular complexes consisting of two identical molecules linked together. - Example aluminium chloride forms a dimer, Al2Cl6. VSEPRT states that the electron pairs around the central atom in a molecule take up positions as far from one another as possible in order to minimise their electrostatic repulsions. The repulsion between lone pairs is greater than the repulsion between lone pairs and bonding pairs which is greater than the repulsion between bonding pairs. -This is used to predict molecule shapes. 51

NES/Chemistry/AS Bonding at AS level The concept of bonding is more complex than the model at IGCSE level. Instead of thinking of bonding as either covalent, or ionic it will be more useful to consider degrees of ionic and covalent character. We use ionic bonding to describe bonding that is more than 50% ionic in character, so the bond pair electrons are more transferred than shared. We use covalent bonding to describe bonding that is more than 50% covalent in character, so the bond pair electrons are more shared than transferred. The degree of bond sharing is dependent on the difference in electronegativity between two bonded atoms. The greater the difference in electronegativity, the more ionic the bond is. 100% covalent bonding is possible as two atoms can have the same electronegativity, but 100% ionic bonding is not possible as there is always some degree of covalent character, even with a large difference in electronegativity. Concepts for AS level Bonding 1. Electronegativity 2. Diagonal relationship 3. Charge density/Proton:electron ratio 4. Strong bonds 5. Weak bonds 6. Electron deficient/Expanded octet 52

NES/Chemistry/AS 1. Electronegativity When two atoms bond together, their electron orbitals overlap and two electrons form a 'bond' between the atoms. This bond is a force of attraction that holds the atoms close together and is how molecules and compounds form. There is an attraction between the shared pair of electrons and the nuclei of both atoms. o x Cl Cl This attractive force for each different atom depends on the nuclear charge, distance and shielding. Electronegativity is the measure of how well an atom can attract electrons from a shared pair. Table of Electronegativities . H 2.1 Li Be B C N O F Ne 1.0 1.5 2.0 2.5 3.0 3.5 4.0 -- Na Mg 0.9 1.2 Al Si P S Cl Ar 1.5 1.8 2.1 2.5 3.0 -- Across a Period, the electronegativity increases because:  the nuclear charge is increasing  the shielding is the same  the distance is the same Down a Group, the electronegativity decreases because:  although the nuclear charge increases  the shielding increases  and the distance increases 53

NES/Chemistry/AS  Example 1: Equal electronegativities making a 100% covalent molecule F o F x  Example 2: Slight difference in electronegativities making a more than 50% covalent molecule δ+ o δ- H x Cl  Example 3: A large difference in electronegativities making a more than 50% ionic character compound. Effectively the sodium atom has lost control of its electron to chlorine and ions form. Na+ o Cl - x 2. Diagonal relationship Moving diagonally on the Periodic Table, you will notice that the electronegativities are similar. Lithium and magnesium for example. This is because the differences in nuclear charge, distance and shielding cancel each other out. This is why the metal/non-metal border on the Periodic Table is drawn on a diagonal line. 3. Charge Density/Proton:Electron Ratio Cations - Charge Density Charge density is a measure of how much electrostatic charge there is in a specific volume, such as inside a a cation. Cations with higher charge densities will share electrons more, so have a greater degree of covalent character.  Across a Period, the charge density of cations increases as the charge is increasing, but the size is similar.  Down a Group, the charge density of cations decreases as the charge is the same and the size of the ion is increasing. 54

NES/Chemistry/AS Anions - Proton:Electron Ratio For anions we focus more on proton:electron ratio instead of charge density. Anions gain electrons, which causes the anion to increase in size. So anions have the same number of protons holding on to more electrons. This means that the nucleus has less control over its own electrons. Anions with a lower proton:electron ratio (greater negative charge) will share electrons more, so have a greater degree of covalent character. Across a Period, the proton:electron ratio decreases and the nucleus exerts less control over its electrons.  Example 4: Beryllium chloride will have a high degree of covalent character (in fact it is over 50% covalent) because of the high charge density of beryllium.  Example 5: Lithium phosphide will have a high degree of covalent character as phosphorus has a low proton:electron ratio. 4. Strong Bonds There are three types of strong bond which act between atoms:  Ionic - between atoms with a large difference in electronegativities  Covalent - between atoms with a low/no difference in electronegativities  Metallic - between metal atoms They are of similar strength. 55

NES/Chemistry/AS 5. Weak Bonds There are four types of weak bond which act between different molecules. They are listed in order, with the strongest first:  Hydrogen bonds - between a hydrogen atom in a molecule that is bonded to a nitrogen, oxygen, or fluorine and a nitrogen, oxygen, or fluorine atom in another molecule  Ion-Dipole interactions - between a polar molecule, like water, and an ion; this is how ionic compounds dissolve in water  Permanent dipole-Permanent dipole - between different molecules with permanent dipoles (polar molecules)  Instantaneous dipole-Induced dipole - between different molecules without dipoles (non-polar molecules) 6. Electron Deficient/Expanded Octet Electron Deficient Elements in Groups 2 and 3 that bond covalently will only make 2, or 3 bonds; so they will only have 4, or 6 electrons in their outer energy level. The energy level is not full, but it is still a stable arrangement. This is different from IGCSE level where you were told that shell had to have 8 electrons in them to be stable (octet rule). See example 10. Expanded Octet Elements in Period 3, or below, that bond covalently will have empty d-orbitals that have yet to be filled. They can use these vacant d-orbitals to form more than four bonds and thus hold more than 8 electrons, e.g. SF6 56

NES/Chemistry/AS 3.1 Ionic Bonding Ionic bonds form between two atoms when the difference in electronegativity is large enough, about 1.7 or greater. The bonding is over 50% ionic in character and the electrons are considered transferred, forming cations and anions.  Example 6: Sodium chloride has an electronegativity difference of 2.1, so the one valance electron from sodium is considered transferred to chlorine's outermost energy level. This forms a sodium ion, Na+ (electronic configuration 1s22s22p6) and a chloride ion, Cl- (electronic configuration 1s22s22p63s23p6). There is an electrostatic force of attraction between the two oppositely charged ions which holds the ions together in a lattice of alternating ions.  Example 7: Dot-cross diagram of calcium fluoride, CaF2 – [Ca]2+ 2x X F  Example 8: Lattice structure of sodium chloride. Cl– surrounded by 6 Na+ ions Na+ surrounded by 6 Cl– ions 57

NES/Chemistry/AS 3.2 Covalent Bonding Covalent bonds form between two atoms when the difference in electronegativity is small, about 1.6 or less. The bond is over 50% covalent in character and the electrons are considered shared, forming molecules.  Example 7: Hydrogen chloride has an electronegativity difference of 0.9, so it forms a covalent molecule, with the bond pair of electrons shared between the hydrogen and the chlorine.  Example 8: Table of dot-cross diagrams for covalent molecules: Name and Formula Stick Diagram Dot-Cross Diagram Hydrogen, H2 H-H HH xx xx x x xx Oxygen, O2 O=O xx xx OO Chlorine, Cl2 Cl - Cl Cl Cl Hydrogen chloride, HCl H - Cl H Cl xx xx xx xx Carbon dioxide, CO2 O=C=O xx Methane, CH4 xx Ethene, C2H4 H OC O H-C-H H H HCH HH H C=C HH H H C C 58 H H

NES/Chemistry/AS Co-ordinate (Dative Covalent) Bond This is a type of covalent bond where one atom gives both electrons to form the covalent bond. Otherwise, it is the same as a regular covalent bond. One atom, or molecule, has a lone pair of electrons that it can give to form the bond, the other atom, or molecule, has to be electron deficient.  Example 9: Ammonium ion xx Ammonia has a lone pair of electrons H NH H A hydrogen cation is electron deficient (in fact it has no electrons at all), H+ A co-ordinate bond forms between the ammonia molecule and the hydrogen cation. H An arrow is used to show the direction the electrons are donated xx H NH H  Example 10: Aluminium chloride dimer, Al2Cl6 Aluminium chloride is a covalent molecule (electronegativity difference 1.5) which is electron deficient. Cl x Al Cl x x Cl When gaseous, it is possible for a co-ordinate bond to form between two AlCl3 molecules. Cl Cl Cl Al Al Cl Cl Cl 59

NES/Chemistry/AS Orbital Overlap In covalent bonding the orbitals (see Topic 2.3) of adjacent atoms are considered to overlap. There are two types of orbital overlap:  σ bonds from end-to-end overlapping  π bonds from side-to-side overlapping σ Bonds All single covalent bonds are σ bonds. These form from the overlapping of s-orbitals and p-orbitals.  Example 11: Hydrogen bonding to hydrogen, H2 s-orbital overlapping with s-orbital  Example 12: Hydrogen bonding to chlorine, HCl s-orbital overlapping with p-orbital  Example 13: Chlorine bonding to chlorine, Cl2 p-orbital overlapping with p-orbital 60

NES/Chemistry/AS Hybridisation Hybridisation happens when s-orbitals and p-orbitals on an atom mix together producing hybrid orbitals. Hybridisation produces orbitals in which the electron density is more concentrated (larger) in one direction. This allows for greater overlap when it bonds with another orbital releasing more energy and producing a stronger bond. All the hybridised bonds around an atom are at identical energy levels and have identical shapes. The s-orbital can mix with either one, two, or three of the p-orbitals to form: Hybridisation Made From Total Orbitals Total Electrons Example 4 Ethyne sp orbital one s-orbital and 2 6 Ethene one p-orbital 8 Ethane sp2 orbital one s-orbital and 3 two p-orbitals sp3 orbital one s-orbital and 4 three p-orbitals 61

NES/Chemistry/AS  Example 14: sp3 hybridisation in methane. Ground state carbon ↑↓ ↑↑ 2s 2px 2py 2pz One 2s electron is promoted up to the 2pz orbital Excited state carbon ↑ ↑↑↑ 2s 2px 2py 2pz All of the orbitals then hybridise sp3 hybrid orbitals ↑ ↑↑ ↑ 2sp3 Four sp3 hybrid orbitals can then overlap with the s-orbitals of four hydrogen atoms to form methane. H s orbital sp3 hybrid orbitals C H H H 62

NES/Chemistry/AS π Bonds A π bond forms after a single bond to make the second bond in a double bond, or the second and third bonds in a triple bond. They form from the sideways overlapping of p-orbitals.  Example 15: sp2 hybridisation and the forming of a double bond in ethene. Ground state carbon ↑↓ ↑↑ 2s 2px 2py 2pz One 2s electron is promoted up to the 2pz orbital Excited state carbon ↑ ↑↑↑ 2s 2px 2py 2pz Three of the orbitals then hybridise, one p-orbital remains unhybridised. sp2 hybrid orbitals ↑↑ ↑ ↑ 2sp2 2p The three sp2 hybrid orbitals overlap with orbitals from carbon and hydrogen to form three σ bonds. H H C C H H 63

NES/Chemistry/AS Then the remaining p-orbital electron on each carbon overlap at 900 to the σ bonds, to form the second bond (π bond) in the double bond.  bonds 2 lobes of the  bond H H C C H H It is a difficult picture to see what is happening. This alternative diagram of ethene might help you visualise it. 64

NES/Chemistry/AS  Example 16: sp hybridisation and the forming of a triple bond in ethyne, C2H2 Ethyne form a carbon-carbon triple bond. This consists of an sp- hybrid σ bond and two sideways p-orbital π bonds. 65

NES/Chemistry/AS Shapes of Covalent Molecules The shapes of molecules and ions depends upon the repulsion between both bonding pairs of electrons and lone pairs of electrons around a central ion, or atom and is predicted using the valence shell electron pair repulsion theory (V.S.E.P.R.T.) For each different combination of lone pairs and bond pairs there is a name and bond angle for that shape. Shape Stick Diagram Bonding Lone Bond Pairs Pairs Angle Linear 2 0 180 (2-dimensional) 120 Trigonal 30 (2-dimensional) Tetrahedral 4 0 109.5 (3-dimensional) Trigonal 90 Bipyramidal 5 0 and (3-dimensional) 120 Octahedral 6 0 90 (3-dimensional) 66

NES/Chemistry/AS As lone pairs of electrons are closer to the central ion, or atom, they repel other electrons slightly more than bond pairs of electrons. Each lone pair of electrons takes approximately 2.5o off the bond angle of the molecule. Shape Stick Diagram Bonding Lone Bond Angle Pairs Pairs (120 - 2.5) = 117.5 Non-linear .. (2-dimensional) 21 Pyramidal .. 3 1 (109.5 - 2.5) = 107 .. .. 2 2 (109.5 - 5.0) = 104.5 (3-dimensional) Non-linear (3-dimensional)  Example 17: Methane, CH4 H Dot-Cross Diagram HCH Stick Diagram H H C H H H Bond Pairs 4 Lone Pairs 0 Bond Angle 109.50 Shape of Molecule Tetrahedral 67

 Example 18: Carbon dioxide, CO2 NES/Chemistry/AS Dot-Cross Diagram Stick Diagram OCO Bond Pairs OCO Lone Pairs 2 (double pairs) Bond Angle 0 Shape of Molecule 1800 Linear  Example 19: Ammonia, NH3 Dot-Cross Diagram HNH H Stick Diagram N H Bond Pairs H Lone Pairs Bond Angle H Shape of Molecule 3 1 109.5 - 2.5 = 1070 Pyramidal 68

NES/Chemistry/AS  Example 20: Boron chloride, BCl3 Cl Dot-Cross Diagram B Cl Cl Note - BCl3 is electron deficient, it forms a stable arrangement without having a full outer energy level. Stick Diagram Cl Cl B Cl Bond Pairs 3 Lone Pairs 0 Bond Angle 1200 Shape of Molecule Trigonal 69

 Example 21: Sulfate ion, SO42- NES/Chemistry/AS Dot-Cross Diagram 2 O OSO O Note - sulfur is in Period 3, so it has vacant d-orbitals and can expand its octet and hold more than 8 electrons in its outer energy level. Note - as sulfate is an ion, it must be put in square brackets and the charge placed in the top right-hand side corner. Note - as a 2- ion, there are 2 extra electrons in the sulfate ion. They are on the oxygen atoms with the single bonds. Stick Diagram O2 S O O O Bond Pairs 4 (2 double, 2 single) Lone Pairs 0 Bond Angle 109.50 Shape of Molecule Tetrahedral 70

NES/Chemistry/AS  Example 22: Nitrate ion, NO3- This is one of the hardest ions - as nitrogen is in Period 2 it cannot expand its octet. Dot-Cross Diagram O N OO Note - there is a co-ordinate bond between one of the oxygens and the nitrogen. Stick Diagram O N Bond Pairs OO Lone Pairs Bond Angle 3 (1 double, 2 single) Shape of Molecule 0 1200 Trigonal 71

NES/Chemistry/AS 3.3 Intermolecular Forces, Electronegativity and Bond Properties Intermolecular Forces There are 2 main types of intermolecular forces (forces of attraction between covalent molecules): 1. Hydrogen bonding 2. van der Waals' a. Permanent dipoles b. Temporary dipoles There are also ion-dipole interactions, but this is not covered until A2 level. It involves the formation of bonds between polar molecules (like water) and ions to break apart ionic lattices - you know this as dissolving. All intermolecular bonds are weak bonds. 1. Hydrogen Bonding Nitrogen, oxygen and fluorine have high electronegativities as well as being small atoms. When nitrogen, oxygen, or fluorine bond to a hydrogen atom (also small) the molecule develops the ability to form hydrogen bonds with other molecules. A hydrogen bond is a strong intermolecular bond (it is still a weak bond when compared to covalent, ionic, or metallic bonds). It is shown as a bar code like dotted line between molecules.  Example 23: Water δ- This means that water, and other molecules that have hydrogen bonding, have higher than expected melting and boiling point for covalent molecules. 72

NES/Chemistry/AS  Example 24: Ammonia dissolved in water δ- Temperature, oCNote - for anything to dissolve in water, it must bond with the water molecules. Graph Showing the Higher Than Expected Boiling Point of Water H2O H2Te H2Se H2S Atomic Number expected boiling point for H2O 73

NES/Chemistry/AS 2. van der Waals' van der Waals' forces of attraction also act between molecules. They are weaker than hydrogen bonds. a. Permanent Dipoles Some molecules have a permanent dipole. A dipole is when opposite ends of a molecule have different partial charges (δ+ and δ-). Dipoles are formed when molecules have polar bonds, due to the difference in electronegativities of the atoms being bonded together.  Example 25: Hydrogen chloride polar bond δ+ x δ- H o Cl When a molecule has a polar bond, it will make the molecule polar as well. However, if a molecule has polar bonds, but the molecule is symmetrical then the molecule will not be polar.  Example 26: Carbon dioxide δ- δ+ δ- OCO The carbon-oxygen double bonds are polar, but the molecule is not polar as the opposite ends are both δ- Carbon dioxide is a non-polar molecule, so it will not form permanent dipole van der Waals' forces of attraction. This type of van der Waals' force of attraction is between a molecule with a permanent dipole and another molecule with a permanent dipole. 74

NES/Chemistry/AS b. Temporary Dipoles Temporary dipole forces are weaker than permanent dipole forces. They occur in molecules that are non-polar, either molecules with non-polar bonds, or symmetrical molecules with polar bonds. See example 26 - carbon dioxide. Even though there are no permanent dipoles in the molecule, the electrons move around randomly inside the molecule. This means that there is a chance that the electrons will be more on one side of the molecule at any given instant. evenly arranged electrons unevenly arranged electrons non-polar polar This polarity is temporary as the electrons will continue to move randomly. However, any temporary polarity will briefly induce polarity in neighbouring molecules. This gives rise to the attractive forces allowing vapour (gas) to condense to liquid. 75

NES/Chemistry/AS Bond Properties There are two bond properties: 1. Bond Energy 2. Bond Length These affect activation energy and therefore stability and reactivity. 1. Bond Energy Bond energy (or bond strength) is the measure of strength of a covalent bond. It is how much energy is required to break a bond, or is released when forming a bond. Common bond energies are given in the data booklet. See also Topic 5.1 b. ii. 2. Bond Length Bond length is the distance between the two nuclei in a covalent bond. The important thing about bond length is its relationship with bond energy. In a covalent bond, the two atoms are held together because both nuclei are attracted to the same pair of electrons. In a longer bond, the shared electron pair is further from at least one of the two nuclei, and so the attractions are weaker. The reactivity of covalent molecules depends on the strength of the bonds being broken (activation energy) as well as the attraction between polar reagents. 76

NES/Chemistry/AS 3.4 Metallic Bonding Metal elements can form metallic bonds. Either atoms of the same element (metal), or a metal and at least one other element (alloy). A metallic bond is a lattice of cations surrounded by a sea of delocalised electrons. Although the structure is made up of cations, each cation has access to the delocalised electrons, so the overall structure is neutral. e e e e e delocalised mobile e e e valence electrons e e that have separated from the atoms e ee e lattice of e metal cations e All delocalised electrons are equally attracted to all the cations, so can move from one cation to another very easily. All localised electrons are limited to existing within the volume of space described by the orbitals of a given atom. 77

NES/Chemistry/AS 3.5 Bonding and Physical Properties Bonding Bond Strength Melting- Electrical Solubility Ionic Boiling Point Conductivity Strong Usually Soluble Compounds High High when Intramolecular - molten, or Insoluble Covalent Strong Low aqueous - Except Molecules Increases with molecules with Intermolecular - Low hydrogen bonding Metallic Weak molecule Insoluble size/length High - Except some Strong reactive metals Usually react with water High Boiling point is proportional to the strength of the weakest type of bonding in a liquid. Melting point is proportional to the strength of the weakest type of bonding in a solid, as well as the packing in the solid lattice. 78

NES/Chemistry/AS Topic 4 - States of Matter The study of the particles in solids, liquids and gases and the interactions between them is important in understanding the physical properties of substances. 4.1 The gaseous state: ideal and real gases and pV = nRT a) state the basic assumptions of the kinetic theory as applied to an ideal gas b) explain qualitatively in terms of intermolecular forces and molecular size: (i) the conditions necessary for a gas to approach ideal behaviour (ii) the limitations of ideality at very high pressures and very low temperatures c) state and use the general gas equation pV = nRT in calculations, including the determination of Mr 4.2 The liquid state a) describe, using a kinetic-molecular model, the liquid state, melting, vaporisation, vapour pressure 4.3 The solid state: lattice structures a) describe, in simple terms, the lattice structure of a crystalline solid which is: (i) ionic, as in sodium chloride, magnesium oxide (ii) simple molecular, as in iodine and the fullerene allotropes of carbon (C60 and nanotubes only) (iii) giant molecular, as in silicon(IV) oxide and the graphite, diamond and graphene allotropes of carbon (iv) hydrogen-bonded, as in ice (v) metallic, as in copper 79

NES/Chemistry/AS b) discuss the finite nature of materials as a resource and the importance of recycling processes c) outline the importance of hydrogen bonding to the physical properties of substances, including ice and water (for example, boiling and melting points, viscosity and surface tension) d) suggest from quoted physical data the type of structure and bonding present in a substance Kinetic Theory is the behaviour of particles in solids, liquids and gases. Ideal Gas is model of gas behaviour that obeys the ideal gas law. Gases are not actually ideal, but if we make certain assumptions then the ideal gas equation gives fairly accurate calculations. Real Gas is how gases actually behave. As conditions change (particularly temperature and pressure) the ideal gas equation no longer gives accurate values. Vapour Pressure is the pressure of an evaporated gas above the surface of the liquid it came from. Crystals are solids, which have ordered structures and symmetry. A Lattice is a regular, repeated three-dimensional arrangement of particles. Lattice is used to describe the structure of crystals. Effectively crystal and Lattice mean the same thing. Viscosity is the magnitude of internal friction in a liquid, think of it as how thick/sticky a liquid is. Surface Tension of a liquid results from an imbalance of intermolecular attractive forces, the cohesive forces between molecules: • A molecule in the bulk liquid experiences cohesive forces with other molecules in all directions • A molecule at the surface of a liquid experiences only net inward cohesive forces 80

NES/Chemistry/AS 4.1 The Gaseous State Ideal Gases The basic assumptions of the kinetic theory as applied to an ideal gas are: 1. The volume of the gas particles is negligible compared with the volume of the container because the gas particles are tiny and spread out 2. All collisions by the gas particles are elastic (no energy is lost) 3. There are no intermolecular forces of attraction between the gas particles 4. The average kinetic energy of the molecules in a sample of gas is directly proportional to the absolute temperature (Kelvin scale) 5. Gas particles move rapidly and randomly in straight lines, colliding occasionally with each other and with the wall of the container For a gas to approach ideal behaviour it needs to:  Have minimal size/Ar  Have minimal intermolecular forces of attraction This occurs at:  Lower size/Ar - weaker intermolecular forces of attraction  Low pressure - the particles are further apart  High temperature - the kinetic energy of the particle is considerably higher than the intermolecular forces of attraction 81

NES/Chemistry/AS Particle size/Ar Note - for an ideal gas, the line is horizontal showing that the volume is constant. Helium is most ideal in its behaviour as it has the smallest size (and lowest Ar) Temperature For an ideal gas, the volume is proportional to the temperature at constant pressure. 82

NES/Chemistry/AS Pressure For an ideal gas, the volume is inversely proportional to pressure at constant temperature. General Gas Equation These assumptions allow us to use the general gas equation: Symbol Quantity Units p Pressure Pa V Volume m3 n Amount (Moles) mol R Gas Constant JK-1mol-1 T Temperature K Note - You can use kPa instead as long as you use dm3 as well. 83

NES/Chemistry/AS Unit Conversions Pressure: Pa kPa atm 100000 100 1 1 atmosphere (atm) is standard atmospheric pressure, equal to 10000 Pascals (Pa) Volume: cm3 dm3 m3 1000000 1000 1 Temperature: 0C K 25 298 The Kelvin scale (or absolute scale) is used where the coldest possible temperature is 0 K. To convert from 0C to K just add 273. Amount: This is just the number of moles of gas. Gas Constant: This is a constant, which is quoted in the data booklet as 8.31 JK-1mol-1 84

NES/Chemistry/AS  Example 1: Calculate the pressure exerted on a 10cm3 sealed container by 0.002 moles of ammonia gas at 25 0C. Volume m3 Volume Temperature K Temperature Pa atm Substituting the mole equation for n we can also get: Symbol Quantity Units m Mass g Mr Relative Molecular Mass gmol-1 85

NES/Chemistry/AS  Example 2: Calculate the Mr of 0.513g of an alkene at 0.1atm, 270C in a 2dm3 sealed container. gmol-1 Real Gases As gases deviate from ideal behaviour they are described as real gases and the general gas equation no longer gives accurate values. As you can see from the graph volume is no longer constant, even though the gases would be in a fixed volume, sealed container. Gases become real, rather than ideal when:  The intermolecular forces are significant - they can no longer be assumed to be negligible  High pressure - the particles take up a significant volume in the container as they are closer together  Low temperature - the kinetic energy of the particles is low enough to be comparable in magnitude to the intermolecular forces of attraction 86

NES/Chemistry/AS 4.2 The Liquid State In a liquid particles are close together, but can slip and slide past each other in continuous motion. Melting Solids have a fixed volume and fixed shape because the attractive forces are strong and hold the particles in a lattice. Energy is required to change a substance from a solid to a liquid and then to a gas. As a solid is heated the particles vibrate more until they have enough energy to overcome the forces of attraction holding them in fixed position in the lattice and start to slip and slide closely over each other. The solid melts. Whilst a pure solid melts the temperature remains constant because the energy supplied is used to overcome the forces of attraction holding the particles in the lattice. When all of the solid has melted, the temperature starts to rise and the particles have more kinetic energy and move faster. Vaporisation This is when a liquid changes into a gas. There are two possible ways of this happening: 1. Boiling 2. Evaporation 1. Boiling At the boiling point the temperature stays constant as the energy supplied is used to overcome the forces of attraction holding particles close together in the liquid state. The entire body of the liquid boils. When all the liquid has boiled to form a gas, the temperature rises again as the particles move faster. 87

NES/Chemistry/AS 2. Evaporation Evaporation is the process when a liquid changes to a gas at temperatures below the boiling point. Only particles at the surface, which happen to have enough energy, can escape from the surface of the liquid to form a gas. Evaporation happens faster as temperature rises. Vapour Pressure At any temperature, the particles in a liquid have a range of kinetic energies. Some particles have enough energy to overcome the forces holding them in the liquid state and they evaporate. Particles with highest kinetic energy evaporate from the liquid surface, so the particles with lower kinetic energies remain in the liquid. This means that the average kinetic energy of the liquid decreases and the liquid temperature falls. If a liquid is placed in a closed vessel, particles will escape from the liquid surface and some particles will rejoin the liquid. When the rate of evaporation is the same as the rate of condensation then the liquid is in equilibrium with its vapour. Liquid ⇌ Vapour Vapour pressure is the pressure due to the vapour above the liquid at equilibrium colliding with the sides of the container. As the temperature increases the average kinetic energy of the particles in the liquid increases. More particles have enough energy to escape into the gas and collide harder and more frequently with the walls of the container so the vapour pressure increases.  When the vapour pressure is less than atmospheric pressure, the liquid will evaporate.  When the vapour pressure is equal to atmospheric pressure, the liquid will boil. 88

NES/Chemistry/AS 4.3 The Solid State In a solid particles are close together, held in a regular, repeating pattern. The particles only vibrate about a fixed point, they do not move. Lattice Structure of Crystalline Solids There are five different types of lattice: 1. Ionic 2. Giant Molecular 3. Simple Molecular 4. Hydrogen-Bonded 5. Metallic In each type the lattice is made up of different types of particles, but overall there is a similar crystalline structure. 1. Ionic The lattice is made up of alternating, oppositely charged ions.  Example 3: Sodium chloride 89

NES/Chemistry/AS 2. Giant Molecular Carbon and silicon can form giant molecular structures as they have a valency of 4. This gives rise to tetrahedral structures (valency 4) and hexagonal layers when one of the valence electrons delocalise through the whole structure.  Example 4: Diamond covalent bond  Example 5: Silicon dioxide 90

NES/Chemistry/AS  Example 6: Graphite  Example 7: Graphene Graphene is a single layer of graphite. 3. Simple Molecular The lattice is made up of small molecules that have formed a solid crystal.  Example 8: Iodine II I I I I I I I I I I I I I I I I I I I I I II I II I II I I I I I I I II 91

NES/Chemistry/AS  Example 9: C60 fullerenes  Example 10: Nanotube fullerenes Nanotubes are made by rolling a layer of graphene into a tube. 92

NES/Chemistry/AS 4. Hydrogen-Bonded The lattice is made up of small molecules that have formed a solid crystal, similar to simple molecular. However, the structure has hydrogen bonds between molecules.  Example 11: Ice When water changes to ice, the hydrogen bonds which hold the molecules together are in fixed positions spacing the molecules further apart than they are in liquid water. 93

NES/Chemistry/AS 5. Metallic The lattice is made up of metal cations, surrounded by a sea of delocalised electrons. There are different types of crystal structure that can be made from this. Properties of Crystal Solids Crystal Hardness Malleability Melting- Electrical Boiling Point Conductivity Ionic Hard Brittle Molten, or Aqueous Giant Molecular Hard - High Simple Molecular Weak - High Poor* Metallic Hard Low Poor Malleable Usually High Good *Graphite and graphene conduct electricity. Hydrogen bonded simple molecules have higher melting and boiling points than simple molecular crystals. 94

NES/Chemistry/AS Properties of Fullerenes Fullerene Hardness Strength Melting- Electrical Reactivity Boiling Conductivity C60 Soft - Point Reactive Nanotube - Strong Poor - Low High Graphene - Very Strong High Very Very High reactive - Resources and Recycling Resources are:  Useful materials extracted from ores/underground  Limited in supply  Expensive to produce  Difficult to dispose of after use  Require energy to produce Resources can be recycled to:  Reduce processing costs  Conserve limited supplies of raw materials  Avoid taking up space in landfill sties  Avoid producing toxic gases when burned However, it can be quite difficult to collect and sort different resources in order to recycle them. 95

NES/Chemistry/AS Topic 5 - Chemical Energetics Enthalpy changes accompany chemical reactions. This topic demonstrates why some reactions and processes are spontaneous and others are not. 5.1 Enthalpy change, ΔH a) explain that chemical reactions are accompanied by energy changes, principally in the form of heat energy; the energy changes can be exothermic (ΔH is negative) or endothermic (ΔH is positive) b) explain and use the terms: (i) enthalpy change of reaction and standard conditions, with particular reference to: formation, combustion, hydration, solution, neutralisation, atomisation (ii) bond energy (ΔH positive, i.e. bond breaking) (iii) --A2 Only-- c) calculate enthalpy changes from appropriate experimental results, including the use of the relationship enthalpy change, ΔH = –mcΔT d) --A2 Only-- 5.2 Hess’ Law a) apply Hess’ Law to construct simple energy cycles, and carry out calculations involving such cycles and relevant energy terms, with particular reference to: (i) determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion (ii) average bond energies (iii) and (iii)--A2 Only-- b) construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the reaction and of the activation energy 96

NES/Chemistry/AS Enthalpy (H) is the energy required to break a bond between gaseous atoms. It is also called bond energy and bond strength. Change in Enthalpy (∆H) is the change in bond energy during a chemical reaction. Exothermic reactions transfer heat energy from the reacting chemicals to their surroundings - they give off heat energy, so ∆H is negative. Endothermic reactions transfer heat from the surroundings to the reacting chemicals - they take in heat energy, so ∆H is positive. Standard Conditions is normal room conditions of 10000Pa pressure, 298K temperature, 1moldm-3 concentration for solutions, Each substance is in its normal physical state. Standard Enthalpy Change of Formation is the enthalpy change when one mole of a compound is formed from its elements under standard conditions. Combustion is the enthalpy change when one mole of a substance is burnt in excess oxygen under standard conditions. Neutralisation is the enthalpy change when one mole of water is formed by the reaction of an acid with an alkali under standard conditions. Atomisation is the enthalpy change when one mole of gaseous atoms is formed from its element under standard conditions. Reaction is the enthalpy change when the amounts of reactants shown in the equation react to give products under standard conditions. Hydration of an anhydrous salt is the enthalpy change when one mole of a hydrated salt is formed from one mole of the anhydrous salt under standard conditions. Solution is the enthalpy change when one mole of solute is dissolved in a solvent to form an infinitely dilute solution under standard conditions. 97

NES/Chemistry/AS Hess' Law states that the total enthalpy change in a chemical reaction is independent of the route by which the chemical reaction takes place as long as the initial and final conditions are the same. Hess' Law Energy Cycle is a diagram to help calculate enthalpy changes. They look like vector triangles. Reaction Pathway Diagram is used to help calculate enthalpy changes. They look like three platforms on a graph. Activation Energy is the minimum energy that colliding particles must possess for a successful collision that results in a chemical reaction. Kinetic Stability - a compound is kinetically stable when the activation energy is high - the reaction is difficult to start. Thermodynamic Stability - a chemical change is thermodynamically stable when ∆H is (a large) positive value - the reaction is difficult to keep going. 98

NES/Chemistry/AS 5.1 Enthalpy Change, ΔH Enthalpy Enthalpy is the energy in a bond between two atoms. Some average bond energies are quoted in the data booklet in kJmol-1. There is no ± sign with these values. Breaking bonds is endothermic (positive ∆H) and bond making is exothermic (negative ∆H). The values are averages as the bond energy does vary depending on what other atoms and bonds are in the chemical. All chemicals are held together by bonds, so each different chemical has its own fixed enthalpy value. To find the enthalpy of a chemical you just have to add up all the bond energies for the chemical bonds in the chemical.  Example 1: What is the enthalpy of water, H2O From the data booklet the value for an O-H bond is 460, so Water enthalpy = 2 x 460 = 920kJmol-1 Note - this type of question has not been asked in the exams, it is given to help you understand the concept of enthalpy. Enthalpy Changes All chemical reactions involve the breaking of bonds in reagents and the forming of new bonds in products. As the enthalpies of products will be different from the enthalpies of the reactants, then there will be a change in enthalpy for every chemical reaction. If the bond is being broken (reagent) the bond energy value from the data booklet is positive - in other words this is how much energy has to be supplied to break the bond. If the bond is being made (product) the bond energy value from the data booklet is negative - in other words this is how much energy is being released as the bond forms. 99