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Williamwood High School CfE Advanced Higher Chemistry Learning Outcomes Name:________________________________ Class: ______ Unit 1 Inorganic and Physical Chemistry Key areas: Electromagnetic radiation and atomic spectra Electromagnetic waves Electromagnetic radiation may be described in terms of waves and can be specified by its wavelength ( ) in metres or an appropriate sub-multiple (e.g. nanometres) and by its frequency (f) in s-1 which is the reciprocal of time in seconds also called the Hertz (Hz). The velocity (c) of electromagnetic radiation is constant and has a value of approximately 3 x 108 ms-1. The electromagnetic spectrum is the range of frequencies or wavelengths of electromagnetic radiation. Velocity, frequency and wavelength are related in the expression: c = f . Dual nature of electromagnetic radiations- waves and particles Under certain circumstances electromagnetic radiation may be regarded as a stream of particles, rather than waves. These particles are known as photons. The energy (E) of radiation, and the energy associated with photons, is related to frequency by Planck’s constant (h) in the expressions: E = h E = L h for one photon for one mole of photons (where L is Avogadro’s constant). Atomic emission The emission spectrum of hydrogen provides evidence of energy levels. Quantum theory states that matter can only emit or absorb energy in small fixed amounts (called quanta). The energy of a bound electron in an atom is quantised. An atom can be considered as emitting a photon of light energy when an electron moves from a higher energy level to a lower energy level Each line of the emission spectrum represents radiation of a specific wavelength or frequency from which the difference in energy between the levels can be calculated. Each element produces a unique pattern of frequencies of radiation in its emission spectrum. Atomic Spectroscopy Atomic emission spectroscopy and atomic absorption spectroscopy are used to identify and quantify the elements present in a sample. Atomic emission spectroscopy and atomic absorption spectroscopy involve transitions between electronic energy levels in atoms. Generally, the energy differences correspond to the visible region of the electromagnetic spectrum, i.e. to the approximate wavelength range of 400-700nm. Some applications use the ultra-violet region (wavelength region approximately 200-400nm). 1 In emission spectroscopy the sample is energised by heat or electricity causing electrons to be promoted to higher energy levels. The wavelength of the radiation emitted as electrons fall back to lower energy levels is measured. In absorption spectroscopy electromagnetic radiation is directed at the atomised sample. Radiation is absorbed as electrons are promoted to higher energy levels. The wavelength of the absorbed radiation is measured. Each element provides a characteristic spectrum which can be used to identify an element. The concentration of an element within a sample can be determined quantitatively if the intensity of emitted or transmitted radiation is measured. Key areas: Atomic orbitals, electronic configurations and the Periodic Table Atomic orbitals Emission spectra of elements with more than one electron provide evidence of sublevels within each principal energy level above the first. Any electron within an atom can be uniquely identified by a set of four quantum numbers: The principal quantum number (n) The angular momentum quantum number (l) The magnetic quantum number (m) The spin magnetic quantum number (s) The principal energy levels correspond to the principal shells (principal quantum number (n) and takes the values 1, 2, 3……). The second and subsequent principal shells contain subshells which correspond to the sublevels (second quantum number (l) defines the subshell type i.e. s, p, d or f). The types of subshell within each principal shell are as follows: First shell s subshell Second shell s and p subshells Third shell s, p and d subshells Fourth shell s, p, d and f subshells An atomic orbital is a region or volume of space within which there is a high (95%) probability of finding an electron. There are four types of orbitals (within subshells), s, p, d and f, each with characteristic shapes. Diagrams of d orbitals can be recognised. Electronic configurations An orbital holds a maximum of two electrons, as required by the Pauli Exclusion Principle which states that no two electrons have the same set of four quantum numbers. The third quantum number (m) determines the actual orbital in which the electron is located. The number of orbitals in each subshell is as follows: s subshell one s orbital p subshell three p orbitals d subshell five d orbitals f subshell seven f orbitals 2 In an isolated atom the orbitals within each subshell are degenerate ( have the same energy) The Aufbau Principle states that orbitals are filled in order of increasing energy. The relative energies corresponding to each orbital can be represented diagrammatically for the first four shells of a multi-electron atom. The fourth quantum number (s) is related to the electron’s spin direction. Hund’s Rule states that when degenerate orbitals are available, electrons fill each singly, keeping their spins parallel before spin pairing starts. Electronic configurations using spectroscopic notation and orbital box notation can be written for elements of atomic numbers 1 to 36. The Periodic Table and ionisation energies The Periodic Table can be subdivided into four blocks (s, p, d and f) corresponding to the outer electronic configurations of the elements within these blocks. The variation in first ionisation energy with increasing atomic number for the first 36 elements can be explained in terms of relative stability of different electron configurations, and so provides evidence for these electronic configurations. The relative values of first, second and subsequent ionisation energies can be explained in terms of the stabilities of the electronic configurations from which the electrons are being removed. Shapes of molecules and polyatomic ions Valence Shell Electron Pair Repulsion Theory (VSEPR) VSEPR theory provides a quick method to predict the shape of a molecule or polyatomic ion from the numbers of bonding electron pairs and non-bonding electron pairs. Arrangement of electron pairs Total number of electron pairs Linear 2 Trigonal planar 3 Tetrahedral 4 Trigonal bipyramidal 5 Octahedral 6 Electron pair repulsions decrease in strength in the order: non-bonding pair/non-bonding pair> non-bonding pair/bonding pair> bonding pair/bonding pair. Differences in strength of electron pair repulsion account for slight deviations from expected bond angles in molecules such as NH3 and H2O. Key area: Transition metals Electronic configurations of transition metal atoms and ions The d-block transition metals are metals with an incomplete d subshell in at least one of their ions. The filling of the d orbitals follows the Aufbau Principle, with the exception of chromium and copper atoms. These exceptions are due to special stability associated with all the d orbitals being half filled or completely filled. When transition metals form ions it is the s electrons which are lost first rather than the d electrons. 3 Oxidation states of transition metals An element is said to be in a particular oxidation state when it has a specific oxidation number. The oxidation number is determined by following certain rules listed on page ___ of notes. Transition metals exhibit variable oxidation states of differing stability. Compounds of the same transition metal but in different oxidation states may have different colours. Oxidation can be considered as an increase in oxidation number and reduction can be considered as a decrease in oxidation number. Compounds containing metals in high oxidation states tend to be oxidising agents whereas compounds with metals in low oxidation states are often reducing agents. Dative covalent bonding of ligands in transition metal complexes A complex consists of a central metal ion surrounded by ligands. Ligands are electron donors and may be negative ions or molecules with non-bonding pairs of electrons. Ligands can be classified as monodentate, bidentate, etc. The total number of bonds from the ligand to the central metal ion is known as the coordination number of the central ion. Complexes are written and named according to IUPAC rules. Colour in transition metal complexes In a complex of a transition metal the d orbitals are no longer degenerate. The energy difference between subsets of d orbitals depends on the position of the ligand in the spectrochemical series. Colours of many transition metal complexes can be explained in terms of d-d transitions. Light is absorbed when electrons in a lower energy d orbital are promoted to a d orbital of higher energy. If light of one colour is absorbed, then the complementary colour will be observed. UV and visible absorption of transition metal complexes The effects of d-d transitions can be studied using ultra-violet and visible absorption spectroscopy. Ultra-violet and visible absorption spectroscopy involves transitions between electronic energy levels in atoms and molecules where the energy difference corresponds to the ultra-violet and visible regions of the electromagnectic spectrum. The wavelength ranges are approximately 200-400nm for ultra-violet and 400-700nm for visible. An ultra-violet/visible spectrometer measures the intensity of radiation transmitted through the sample and compares this with the intensity of incident radiation. Catalysis by transition metals The presence of unpaired d electrons or unfilled d orbitals allows intermediate complexes to form, providing reaction pathways with lower activation energies compared to the uncatalysed reaction. This allows transition metals or their compounds act as catalysts in many chemical reactions. It is believed that the presence of unpaired d electrons or unfilled d orbitals allows intermediate complexes to form, providing reaction pathways of lower energy compared to the uncatalysed reaction. 4 The variability of oxidation states of transition metals is also an important factor. Homogeneous and heterogeneous catalysts should be explained in terms of changing oxidation states. Key area: Chemical equilibrium Equilibrium expressions, factors affecting equilibrium and calculation of the composition of equilibrium mixtures A chemical reaction is in equilibrium when the composition of the reactants and products remains constant indefinitely. The equilibrium constant (K) characterises the equilibrium composition of the reaction mixture. The equilibrium constant can be measured in terms of concentrations or, for gaseous reactions, in terms of pressure. For the general reaction: aA + bB ⇌ cC + dD The equilibrium constant, K = [C]c [D]d [A]a [B]b Where [A], [B], [C] and [D]are the equilibrium concentrations of A, B, C and D, respectively, and a, b, c and d are the stoichiometric coefficients in a balanced reaction equation. In a homogeneous equilibrium all the species are in the same phase. In a heterogeneous equilibrium the species are in more than one phase. The concentrations of pure solids or pure liquids are constant and are given the value 1 in the equilibrium equation. Equilibrium constants are independent of the particular concentrations or pressures of species in a given reaction. Equilibrium constants depend on the reaction temperature. For endothermic reactions a rise in temperature causes an increase in K, i.e., the yield of the product is increased. For exothermic reactions a rise in temperature causes a decrease in K, i.e., the yield of the product is decreased. The presence of a catalyst does not affect the value of the equilibrium constant. The equilibrium constants have no units. The effects of changes in concentration or pressure on the position of equilibrium can be explained quantitatively in terms of a fixed equilibrium constant. Ionic product of water and calculations of pH The ionisation of water can be represented by: H2O(l) + H2O(l) ⇌ H3O+ (aq) + OH - (aq) Water is amphoteric. 5 The dissociation constant for the ionisation of water is known as the ionic product and is represented by: Kw = [H3O+][OH -] The value of the ionic product varies with temperature. At 25o C the value of Kw is approximately 1 x 10-14 mol2 l-2 . A shorthand representation of H3O+ is H+. Stoichiometric equations and equilibrium expressions can be written using H+ instead of H3O+ where the meaning is clear. The relationship between pH and the hydrogen ion concentration is given by pH = -log10 [H+] and conversely [H+] = 10– pH. Bronsted-Lowry acids and bases The Bronsted–Lowry definitions of acid and base state that an acid is a proton donor and a base is a proton acceptor. For every acid there is a conjugate base, formed by loss of a proton. For every base there is a conjugate acid, formed by gain of a proton. A soluble salt of a strong acid and a strong base dissolves in water to produce a neutral solution. A soluble salt of a weak acid and a strong base dissolves in water to produce an alkaline solution. A soluble salt of a strong acid and a weak base dissolves in water to produce an acidic solution. Soaps are salts of weak acids and strong bases. The pH of a salt solution can be explained by reference to the equilibria between the acids and alkalis involved. Strong and weak acids and bases In aqueous solution, strong acids/bases are completely dissociated into ions but weak acids/bases are only partially dissociated. Examples of strong acids include hydrochloric acid, sulfuric acid and nitric acid. Ethanoic, carbonic and sulfurous acids are examples of weak acids. Solutions of metal hydroxides are strong bases. Ammonia and amines are examples of weak bases. The weakly acidic nature of solutions of carboxylic acids, sulfur dioxide and carbon dioxide can be explained by reference to equations showing the equilibria. The weakly alkaline nature of a solution of ammonia or amines can be explained by reference to an equation showing the equilibrium. The acid dissociation constant is represented by Ka or by pKa where pKa = –log10Ka The dissociation in aqueous solution of an acid of general formula HA can be represented as: HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq) 6 The acid dissociation constant of acid HA is given by: [H 3 O ][A ] Ka [HA] The conjugate base of an acid of general formula HA is A-. The dissociation constant of an acid can be represented by pKa where pKa = - log10Ka The approximate pH of a weak acid can be calculated from the dissociation constant: pH = ½ pKa – ½ log10c The dissociation in aqueous solution of base of general formula B can be represented as: BH+(aq) + OH-(aq) B(aq) + H2O(l) The conjugate acid of a base of general formula B is BH+. The dissociation of the conjugate acid of the base can be represented as BH+(aq) + H2O(l) B(aq) + H3O+(aq) The dissociation constant for the conjugate acid is Ka = B H3O+ BH+(aq) Buffer solutions A buffer solution is one in which the pH remains approximately constant when small amounts of acid or base are added. An acid buffer consists of a solution a weak acid and one of its salts. In an acid buffer solution the weak acid can supply hydrogen ions when these are removed by the addition of a small amount of base. The salt of the weak acid provides the conjugate base, which can absorb excess hydrogen ions produced by the addition of a small amount of acid. A basic buffer consists of a solution of a weak base and one of its salts. In a basic buffer solution the weak base removes excess hydrogen ions and the conjugate acid provided by the salt supplies hydrogen ions when these are removed. The pH of an acid buffer solution can be calculated from its composition and from the acid dissociation constant: pH = pKa –log10 [acid] [salt] Indicators Indicators are weak acids for which the dissociation can be represented as HIn(aq) + H2O(l) H3O-(aq) 7 + In-(aq) The acid dissociation constant is represented as KIn and is given by the following expression: [H 3O ] [In ] K In [HIn] In aqueous solution the colour of the acid is distinctly different from that of its conjugate base. The colour of the indicator is determined by the ratio of HIn to In-. The theoretical point at which colour changes occurs when H+ = KIn The colour change is assumed to be distinguishable when HIn and In- differ by a factor of 10. The pH range over which a colour change occurs can be estimated by the expression pH = pKIn 1 Key area: Reaction feasibility Standard enthalpy of formation The standard enthalpy of formation of a substance (Hof) is the enthalpy change when one mole of a substance is formed from its elements in their standard states. The standard enthalpy of a reaction can be calculated from tabulated standard molar enthalpies of formation using the relationship Hof = Hof(products) - Hof(reactants) Knowledge of standard conditions is important. Entropy The entropy (S) of a system is the degree of disorder of the system. The greater the disorder, the greater the entropy. Entropy increases as temperature increases. Changes of state involve changes in entropy. Melting and evaporation are accompanied by increases in entropy. One version of the Third Law of Thermodynamics states that the entropy of a perfect crystal at 0 K is zero. The standard entropy of a substance is the entropy value for the standard state of the substance. The change in standard entropy for a reaction system can be calculated from the standard entropy values for the reactants and products: So = So(products) - So(reactants) 8 Free energy One version of the Second Law of Thermodynamics states that the total entropy of a reaction system and its surroundings always increases for a spontaneous process. Heat energy released by the reaction system into the surroundings increases the entropy of the surroundings, whereas heat absorbed by the reaction system from the surroundings decreases the entropy of the surroundings. The change in entropy of the surroundings that occurs as a result of a chemical reaction can be calculated from the temperature and from the enthalpy change for the reaction system. The total entropy change is normally expressed in terms of Gibbs free energy (G). The direction of spontaneous change is in the direction of decreasing free energy. The change in standard free energy for a reaction is related to the standard enthalpy and entropy changes by: Go = Ho - T So The standard free energy change of a reaction can be calculated from the standard enthalpy and standard entropy changes for the reaction. The standard free energy change of a reaction can be calculated from the standard free energies of formation of the reactants and products using the relationship: Go = Go(products) - Go(reactants) Applications of the concept of free energy The feasibility of a chemical reaction under standard conditions can be predicted from the calculated value of the change in standard free energy (Go). The temperature at which the reaction becomes feasible can be calculated for a reaction for which both Ho and So have positive values. A reaction is feasible under standard conditions if the change in standard free energy between reactants and products is negative. This means that the equilibrium composition favours the products over the reactants. Under non-standard conditions any reaction is feasible if G is negative. At equilibrium G = 0. A reaction will proceed spontaneously in the forward direction until the composition is reached where G = 0 Key area: Kinetics Determining the order of a reaction The rate of a chemical reaction normally depends on the concentrations of the reactants. 9 For a first order reaction the rate of reaction is proportional to the concentration of one reactant and the rate can be expressed as: rate = k A Where k is the rate constant and A is the concentration of reactant A in mol l-1. The order of a reaction with respect to any one reactant is the power to which the concentration of that reactant is raised in the rate equation. The overall order of a reaction is the sum of the powers to which the concentrations of the reactants are raised in the rate equation. The order of a reaction can only be determined from experimental data. In general for a reaction of type nA + mB products Where the rate expression is of the form: rate = k An Bm The order of reaction is n with respect to A and m with respect to B and the overall order is n+ m. The rate constant can be determined from initial rate data for a series of reactions in which the initial concentrations of reactants are varied. Reaction mechanisms Reaction mechanisms usually occur by a series of steps. The rate of reaction is dependent on the slowest step which is called the ‘rate determining step’. Experimentally determined rate equations can provide evidence for a proposed reaction mechanism but cannot provide proof as other possible reaction mechanisms may also give the same rate equation. 10 Unit 2 Organic Chemistry and Instrumental Analysis Key area: Molecular orbitals Formation of bonding molecular orbitals When atoms approach each other, their separate sets of atomic orbitals merge to form a single set of molecular orbitals. Some of the molecular orbitals, known as ‘bonding molecular orbitals’, occupy the region between two nuclei. The attraction of positive nuclei to negative electrons occupying bonding molecular orbitals is the basis of bonding between atoms. Each molecular orbital can accommodate a maximum of two electrons. Hybridisation Hybridisation is the process of mixing atomic orbitals within an atom to generate a set of new atomic orbitals called hybrid orbitals. Bonding in alkanes can be described in terms of sp3 hybridisation and sigma bonds. Bonding in alkenes can be described in terms of sp2 hybridisation and both sigma and pi bonds. Bonding in alkynes can be described in terms of sp hybridisation with both sigma and pi bonds. A sigma bond () is a covalent bond formed by end-on overlap of two atomic orbitals lying along the axis of the plane. A pi bond is a covalent bond formed by the sideways overlap of two parallel atomic orbitals lying perpendicular to the axis of the bond. The bonding continuum Van der Waal’s forces of attractions describe all intermolecular forces of attraction including London dispersion forces, permanent dipole to permanent dipole and hydrogen bonds. London dispersion forces are weak intermolecular forces of attraction that arise due to a temporary uneven distribution of electrons within atoms and molecules. In a non-polar covalent bond, the bonding molecular orbital is symmetrical about the midpoint between two atoms. Polar covalent bonds result from bonding molecular orbitals which are asymmetric about the midpoint between two atoms. Ionic compounds represent an extreme case of asymmetry with the bonding molecular orbitals being almost entirely located around just one atom. Absorption of visible light by organic molecules Most organic molecules appear colourless because the energy difference between the highest occupied molecular orbital (HOMO) and the lowest unoccupied molecular orbital 11 (LUMO) is relatively large resulting in the absorption of light in the ultraviolet region of the spectrum. Coloured organic compounds contain delocalised electrons within molecular orbitals which extend across several atoms. This is known as a conjugated system. A conjugated system exists when a molecule has alternating single and double bonds. The more atoms the delocalised molecular orbital spans, the smaller the energy gap between the delocalised orbital and the next unoccupied orbital and hence the lower the frequency of light (or longer the wavelength or lower the energy of radiation) absorbed by the compound. When the wavelength of the light absorbed is in the visible region, the organic substance will appear coloured. Molecules in which the structural formula contains alternate double bonds will exhibit molecular orbitals containing delocalised electrons which will extend the conjugated section of the molecule. Chromophores The chromophore is the group of atoms within a molecule which is responsible for the absorption of light in the visible region of the spectrum (i.e. responsible for its colour). This is usually a reasonably large part of the molecule - usually involving a conjugated system. Light can be absorbed when electrons in a chromophore are promoted from one molecular orbital to another. If the chromophore absorbs light of one colour, the compound will exhibit the complementary colour. Key area: Molecular structure Structural formulae Structural formulae and skeletal formulae can be drawn for organic molecules with no more than 10 carbon atoms in their longest chain. Conversion between molecular, structural and skeletal organic molecules can be demonstrated. In a skeletal formula neither the carbon atoms, nor any hydrogens attached to the carbon atoms, are shown. The presence of a carbon atom is implied by a ‘kink’ in the carbon backbone, and at the end of a line. Key area: Stereochemistry Stereoisomerism Stereoisomers have identical molecular formulae and the atoms are bonded together in the same order but the arrangement of the atoms in space is different, making them nonsuperimposable. 12 Geometric isomerism Geometric isomerism is one type of stereoisomerism and can arise due to the lack of free rotation around a bond, frequently a carbon–carbon double bond. Geometric isomers are labelled cis and trans according to whether the substituent groups are on the same side or on different sides of the carbon-carbon double bond. Geometric isomers display differences in some physical properties. Geometric isomerism can also influence chemical properties, for example cis-but-2enedioic acid is more readily dehydrated than trans-but-2-enedioic acid. Optical isomerism Optical isomers are non-superimposable mirror images of asymmetric molecules and are said to be chiral molecules or enantiomers. Optical isomerism can occur in substances in which four different groups are arranged around a carbon atom. Optical isomers have identical physical and chemical properties, except when they are in a chiral environment, but they have an opposite effect on plane polarised light and are said to be optically active. Racemic mixtures contain equal amounts of both optical isomers and are optically inactive. In biological systems only one optical isomer of each organic compound is usually present. Key area: Synthesis Reaction types Equations can be written for the following reaction types and, given equations, these reaction types can be identified. substitution addition elimination condensation hydrolysis oxidation reduction dehydration Synthetic routes Synthetic routes, with no more than three steps, can be devised from a reactant to a product. Molecular structures can be examined and possible reactions can be determined. 13 Bond fission When an organic reaction takes place, bonds are broken and formed. If, when the bond between atoms breaks, each atom retains one electron from the former covalent bond, then two free radicals are formed. This is known as homolytic fission. Reactions involving free radicals tend to result in formation of very complex mixtures of products, thus making them unsuitable for synthesis. If, when the bond between atoms breaks, one atom retains both of the electrons from the former covalent bond, then an ion pair is formed. This is known as heterolytic fission. Reactions proceeding via heterolytic fission tend to produce far fewer products and are therefore better suited for synthesis. Heterolytic fission will be favoured when the bond between the atoms is polar. Electrophiles and nucleophiles In reactions involving heterolytic bond fission, attacking groups are classified as 'nucleophiles' or 'electrophiles'. Nucleophiles are atoms or groups of atoms which are attracted towards atoms bearing a (partial) positive charge. Nucleophiles are capable of donating and sharing an electron pair to form a new bond. Electrophiles are atoms or groups of atoms which are attracted towards atoms bearing a (partial) negative charge. Electrophiles are capable of accepting an electron pair. Curly arrow notation Double-headed and single headed curly arrows are used to show the movement of electron during a reaction. The base of the arrow shows the original location of the pair of electrons. The head of the arrow indicates the destination of the pair of electrons. An arrow starting at the middle of a covalent bond indicates that heterolytic bond fission is occurring. When an arrow is drawn with the head pointing to the space between two atoms, this indicates that a covalent bond will be formed between the two atoms. A double-headed arrow indicates the movement of an electron pair and a single-headed arrow indicates the movement of a single electron. Halogenoalkanes Halogenoalkanes are named according to IUPAC rules. Monohalogenoalkanes can be classified as primary, secondary or tertiary. Monohalogenoalkanes undergo nucleophilic substitution reactions. Monohalogenoalkanes undergo elimination reactions to form alkenes. 14 Monohalogenoalkanes react with: i. alkalis to form alcohols ii. alcoholic alkoxides to form ethers iii. ethanolic cyanide to form nitriles which can be hydrolysed to carboxylic acids (chain length increased by one carbon atom) iv. ammonia to form amines via alkyl ammonium salts. Monohaloalkanes can also undergo elimation reactions to form alkenes. Alkenes Alkenes can be prepared in the laboratory by i. dehydration of alcohols using aluminium oxide, conc. sulfuric acid or orthophosphoric acid. ii. base-induced elimination of hydrogen halides from monohalogenoalkanes. Alkenes undergo i. catalytic addition with hydrogen to form alkanes ii. addition with halogens to form dihalogenoalkanes iii. addition with hydrogen halides according to Markownikoff’s rule to form monohalogenoalkanes iv. acid-catalysed addition with water according to Markownikoff’s rule to form alcohols. The mechanisms of the above electrophilic addition reactions involve i. ii. iii. for halogenation, cyclic ion intermediate for hydrohalogenation, carbocation intermediate for acid catalysed hydration, carbocation intermediate. All of the above mechanisms can be represented using curly arrows. Reaction mechanisms The following reaction mechanisms can be described in terms of electron shifts. radical substitution of alkanes. electrophilic addition to alkenes carbocation mechanism cyclic ion intermediate mechanism. nucleophilic substitution: SN1 and SN2 The dominance of an SN1 or SN2 mechanism for a particular haloalkane can be explained in terms of steric hindrance and the inductive stabilisation of an intermediate carbocation. An SN2 reaction proceeds via a single five-centred transition state, whereas an SN1 reaction occurs in two steps via a carbocation. 15 Physical Properties The following physical properties are explained in terms of the intermolecular forces involved melting and boiling points miscibility with water polarity Alcohols Alcohols exhibit hydrogen bonding and as a result have higher boiling points than other organic compounds of comparable relative formula mass and shape. The lower alcohols are miscible with water but as their chain length increases their solubility in water decreases. Alcohols can be prepared from i. alkenes by hydration. ii. halogenoalkanes by substitution. In industry, alcohols (except methanol) can be manufactured by the acid-catalysed hydration of alkenes. Alcohols react with some reactive metals to form alkoxides. Alcohols can be dehydrated to alkenes. Alcohols undergo condensation reactions slowly with carboxylic acids and more vigorously with acid chlorides to form esters. Alcohols undergo oxidation with oxidizing agents. i Primary alcohols are oxidized to aldehydes and then to carboxylic acids. ii Secondary alcohols are oxidized to ketones. iii Tertiary alcohols are not readily oxidized. Ethers Ethers have the general formula R1-O-R2 where R1 and R2 are alkyl groups. Ethers are named according to IUPAC rules. Due to the lack of hydrogen bonding, ethers have lower boiling points than the corresponding isomeric alcohols. Ether molecules can hydrogen bond with water molecules thus explaining the solubility in water of some ethers of low relative formula mass. The solubility decreases as the molecular size increases. Ethers are highly flammable and on exposure to air may form explosive peroxides. Ethers can be prepared by the reaction of halogenoalkanes with alkoxides. Ethers are useful solvents for many organic compounds due to relative stability with respect to oxidation and reduction. 16 Aldehydes, ketones and carboxylic acids The following physical properties can be explained in terms of dipole-dipole attractions and/or hydrogen bonding i. ii. iii. Tollens’ reagent or Benedict’s solution can be used to distinguish between aldehydes and ketones. Aldehydes reduce the complexed silver(I) ion and the complexed copper(II) ion to silver and copper(I), respectively. Aldehydes and ketones can be reduced to primary and secondary alcohols, respectively, by reaction with lithium aluminium hydride in ether. The reaction with 2,4-dinitrophenylhydrazine to form a 2,4-dinitrophenylhydrazone is an example of a condensation reaction. The melting points of the resulting 2,4-dinitrophenylhydrazones are used to identify carbonyl compounds. Aldehydes are generally more reactive than ketones because the presence of two alkyl groups in ketones hinders nucleophilic attack and reduces the partial positive charge on the carbonyl carbon atom. In pure carboxylic acids hydrogen bonding produces dimers thus explaining the relatively high boiling points. Dimerisation does not occur in aqueous solution. Carboxylic acid molecules also form hydrogen bonds with water molecules thus explaining the appreciable solubility of the lower carboxylic acids in water. As the chain length increases water solubility decreases. Carboxylic acids are weak acids. Their slight dissociation in water can be explained by the stability of the carboxylate ion caused by electron delocalisation. Carboxylic acids can be prepared by: i. ii. Reactions of carboxylic acids include: i. ii. iii. iv higher boiling points than corresponding alkanes. lower boiling points than corresponding alcohols. miscibility of lower members with water. oxidising primary alcohols and aldehydes. hydrolysing nitriles, esters or amides. formation of salts by reactions with metals, carbonates and alkalis. condensation reactions with alcohols to form esters. reaction with ammonia or amines and subsequent heating of the ammonium salt to form amides. reduction with lithium aluminium hydride to form primary alcohols. Amines Amines are named according to IUPAC rules. Amines are organic derivatives of ammonia and can be classified as primary, secondary or tertiary. 17 Primary and secondary amines, but not tertiary amines, associate by hydrogen bonding and as a result have higher boiling points than isomeric tertiary amines and alkanes with comparable relative formula masses. Amine molecules can hydrogen bond with water molecules thus explaining the appreciable solubility of the lower amines in water. The nitrogen atom in amines has a lone pair of electrons which can accept a proton from water, producing hydroxide ions. Amines are weak bases. Amines react with aqueous mineral or carboxylic acids to form salts. Aromatics Bonding in benzene can be described in terms of sp2 hybridisation, sigma and pi bonds and electron delocalisation. Benzene C6H6 is the simplest aromatic hydrocarbon and its unexpected stability can be attributed to the presence of delocalised electrons. A benzene ring in which one hydrogen atom has been substituted by another group is known as the phenyl group and has the formula -C6H5. Most reactions of benzene involve attack of an electrophile on the cloud of delocalised electrons. This is called electrophilic substitution. Benzene resists addition reactions but undergoes electrophilic substitution reactions. These include: i. halogenation e.g. chlorination and bromination to produce chlorobenzene and bromobenzene ii. nitration to produce nitrobenzene iii. sulfonation to produce benzene sulfonic acid iv. alkylation to produce alkylbenzenes. Key area: Experimental determination of structure In organic chemistry a number of experimental techniques are carried out to verify that the correct chemical structure has been synthesised. Elemental microanalysis Elemental microanalysis (combustion analysis) can be used to determine the masses of C, H, O, S and N in a sample of an organic compound in order to find the empirical formula. Other elements in the organic compounds have to be determined by other methods 18 Mass spectrometry Mass spectrometry can be used to determine the accurate formula mass and structural features of an organic compound. A conventional mass spectrometer functions in the following manner. The sample is firstly vaporised and then ionised by being bombarded with electrons. Fragmentation can occur when the energy available is greater than the molecular ionisation energy. The parent ion and ion fragments are accelerated by an electric field and then deflected by a magnetic field. The strength of the magnetic field is varied to enable the ions of all the different mass/charge ratios to be detected in turn. A mass spectrum is obtained. Organic compounds can be identified from the very accurate determination of the relative molecular masses of the parent ion and the ion fragments. Infra red spectroscopy Infra-red spectroscopy can be used to identify certain functional groups in an organic compound. The vibrational energy of a chemical bond is quantised. Infra-red radiation causes parts of a molecule to vibrate. The wavelengths which are absorbed and cause the vibrations will depend on the type of chemical bond and the groups or atoms at the ends of these bonds. Infra-red radiation is passed through a sample of the organic compound and then to a detector which measures the intensity of the transmitted radiation at different wavelengths. Wavelength is expressed in spectra as wavenumber. The unit of measurement of wavenumber is the reciprocal of length in centimetres (cm-1). Proton Nuclear magnetic resonance spectroscopy (1 H NMR) Proton nuclear magnetic resonance spectroscopy (proton NMR) can give information about i. the different environments of hydrogen atoms in an organic molecule ii. how many hydrogen atoms there are in each of these environments Hydrogen nuclei behave like tiny magnets and in a strong magnetic field some are aligned with the field (lower energy) while the rest are aligned against it (higher energy).Absorption of radiation in the radiofrequency region of the electromagnetic spectrum will cause the hydrogen nuclei to ‘flip’ from the lower energy alignment to the higher one. As they fall back from the higher to the lower level the emitted radiation is detected. In the NMR spectrum the peak position is related to the environment of the proton. The area under the peak is related to the number of protons in that environment. 19 An interaction with H atoms on a neighbouring carbon atom can result in the splitting of NMR peaks into 'multiplets'. The number of H atoms on the neighbouring carbon will determine the number of lines within a multiplet. Hydrogen nuclei behave like tiny magnets and in a strong magnetic field some are aligned with the field (lower energy) whilst the rest are aligned against it (higher energy). Absorption of radiation in the radio-frequency region of the electromagnetic spectrum will cause the hydrogen nuclei to ‘flip’ from the lower to the higher energy alignment. As they fall back from the higher to the lower level, the emitted radiation is detected. The standard reference substance used in NMR spectroscopy is tetramethylsilane (TMS) which is assigned a chemical shift value equal to zero. Low resolution proton NMR spectra can be drawn and analysed. High resolution proton NMR spectra can be analysed. Key area: Pharmaceutical chemistry Effect of drugs on the body Drugs are substances which alter the biochemical processes in the body and those which have a beneficial effect are called medicines. A medicine usually contains the drug plus other ingredients. Classification of drugs Many medicines can be classified as agonists or as antagonists according to whether they enhance or block the body’s natural responses. An agonist will produce a response like the body’s natural active compound. An antagonist produces no response but prevents the action of the body’s natural active compound. How drugs work Most medicines work by binding to receptors. Receptors are usually protein molecules that are either on the surface of cells where they interact with small biologically active molecules or are enzymes that catalyse chemical reactions (catalytic receptors). The structural fragment of a drug molecule which confers pharmacological activity upon it normally consists of different functional groups correctly orientated with respect to each other. The overall shape and size of the drug has to be such that it fits a binding site. The functional groups on both the drugs and the receptor are positioned such that the drugs can interact with and bind to the receptor. The types of interaction between drugs and binding sites can be identified. By comparing the structures of drugs that have similar effects on the body, the structural fragment that is involved in the drug action can be identified. 20 Many drugs act as enzyme inhibitors by binding to the enzyme’s active site and blocking the reaction normally catalysed there. % solution is the mass of solute made up to 100 cm3 of solution. % by volume is the number of cm3 of solute made up to 100cm3 of solution. The unit ppm stands for parts per million and refers to 1 mg per kg. 21 Unit 3 Researching Chemistry Overarching principles Throughout each of the key areas the learners must be able to apply the following principles: Precision Accuracy Uncertainties Units Key area: Stoichiometric calculations Stoichiometry is the study of quantitative relationships involved in chemical reactions. The ability to balance and interpret equations enabling calculations to be carried out involving any of the above skills/techniques is an important part of chemistry at this level and is examinable in the course assessment. A quantitative reaction is one in which the substances react completely according to the mole ratios given by the balanced (stoichiometric) equation. Experimental determination of percentage yield Theoretical percentage yields can be calculated and compared with actual yield . Mass transfer or mechanical losses, purification of product, side reactions, equilibrium position and purity of reactants, inevitably lowers the percentage yield. Key area: Gravimetric analysis In gravimetric analysis the mass of an element or compound present in a substance is determined by chemically changing that substance into some other substance of known chemical composition, which can be readily isolated, purified and weighed. Key techniques to perform a gravimetric analysis include ‘weighing by difference’ and ‘heating to constant mass’. Key area: Volumetric analysis Volumetric analysis involves using a solution of accurately known concentration in a quantitative reaction to determine the concentration of another substance. 22 A solution of accurately known concentration is known as a standard solution. A standard solution can be prepared directly from a primary standard. A primary standard must have, at least, the following characteristics: high state of purity stability in air and in solution solubility reasonably high formula mass. The volume of reactant solution required to just complete the reaction is determined by titration. The equivalence point is the point at which the reaction is just complete. The ‘end point’ is the point at which a change is observed and is associated with the equivalence point. An indicator is a substance which changes colour at the end-point. Acid/base titrations are based on neutralisation reactions. Controls Use of a control validates a technique and may consist of carrying out a determination on a solution of known concentration. Complexometric titrations Complexometric titrations are based on complex formation reactions. EDTA is an important complexometric reagent and can be used to determine the concentration of metal ions such as nickel(II). Redox titrations are based on redox reactions. Substances such as potassium manganate (VII) which can act as their own indicators are very useful reactants in redox titrations. Back titrations Back titration is used to find the number of moles of a substance by reacting it with an excess volume of reactant of known concentration. The resulting mixture is then titrated to work out the number of moles of the reactant in excess. From the initial number of moles of that reactant the number of moles used in the reaction can be determined, making it possible to work back to calculate the initial number of moles of the substance under test. A back titration is useful when trying to work out the quantity of substance in an insoluble solid. Key area: Practical skills and techniques Standard solutions and calibration curves Preparation of standard solutions using accurate an accurate dilution technique. 23 Colorimetry uses the relationship between colour intensity of a solution and the concentration of the coloured species present. Calibration curves can be constructed to determine an unknown using solution of appropriate concentration. The concentration of the ‘unknown’ solution is determined from its absorbance and by referring to the calibration curve. The straight line section of the calibration graph should cover the dilution range likely to be used in the determination. Preparation and purification of an experimental product Distillation, reflux, vacuum, filtration, recrystallization and separating funnels are practical skills that can be used in the preparation and purification of an experimental product. Distillation: The boiling point of a compound — determined by distillation — is welldefined and thus is one of the physical properties of a compound by which it is identified. Refluxing: Used to apply heat energy to a chemical reaction mixture over an extended period of time. The liquid reaction mixture is placed in a round-bottomed flask along with anti-bumping granules with a condenser connected at the top. The flask is heated vigorously over the course of the chemical reaction; any vapours given off are immediately returned to the reaction vessel as liquids when they reach the condenser. Vacuum filtration: Using a Buchner, Hirsch or sintered glass funnel. These methods are carried out under reduced pressure and provide a faster means of separating the precipitate from the filtrate. The choice of filtering medium depends on the quantity and nature of the precipitate. Recrystallisation: Lab technique used to purify solids, based upon solubility. The solvent for recrystallisation must be carefully selected such that the impure compound is insoluble at lower temperatures, yet completely soluble at higher temperatures. The impure compound is dissolved gently in the minimum volume of hot solvent then filtered to remove insoluble impurities. The filtrate is allowed to cool slowly to force crystallisation. The more soluble impurities are left behind in the solvent. Use of a separating funnel: Solvent extraction can be an application of the partition of a solute between two liquids. It is based on the relative solubility of a compound in two different immiscible liquids, usually water and an organic solvent. The equilibrium constant is called the partition coefficient. The two solvents form two separate layers in the separating funnel and the lower layer is run off into one container and the upper layer is poured out into another container. The quantity of solute extracted depends on the partition coefficient and on the number of times that the process is repeated. Thin-layer chromatography Thin –layer chromatography (TLC) uses a fine film of silica or aluminium oxide spread over glass or plastic to assess the purity of experimental products. Rf values can be calculated and under similar conditions a compound will always have the same Rf value within experimental error. Since a pure substance will show up as only one spot on the developed chromatogram, TLC can be used to assess the purity of a product prepared in the lab. 24 Chromatographic separations depend on the partition equilibrium between two phases; one stationary and the other mobile. Other examples of chromatography are: paper chromatography gas liquid chromatography In paper chromatography, the stationary phase is the water held on the paper and the mobile phase is another solvent. In gas-liquid chromatography the stationary phase is a liquid held on a solid support and the mobile phase is a gas. Determination of melting point and mixed melting point The melting point of an organic compound is one of several physical properties by which it can be identified. A crystalline substance has a sharp melting point falling within a very small temperature range. Determination of the melting point can also give an indication of the purity of an organic compound, as the presence of impurities lowers the melting point and extends its melting temperature range. Since impurities lower the melting point, the technique of mixed melting point determination can be used as a means of identifying the product of a reaction 25