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Ministry of Health of Russian Federation
Federal State Autonomous Educational Institution of Higher Education “N.I. Pirogov Russian National Research Medical University”
Ministry of Health of Russian Federation
FGAOU RNIMU named after N.I. Pirogov of the Ministry of Health of Russia
The program of the Entry Test in Chemistry
I. Scope of application and regulatory references
The program of the entry test was developed for applicants to the Federal State Autonomous Educational Institution of Higher Education N.I. Pirogov RNIMU of the Ministry of Health of Russia for higher education programs: bachelor’s and specialist’s programs, based on the requirements of federal state educational standards for secondary vocational education in specialties related to the field of knowledge “Health and medical sciences”, the requirements of federal state educational standards for specialties related to areas of knowledge “Health and medical sciences”, “Education and pedagogical sciences”, “Mathematical and natural sciences”, “Agriculture and agricultural sciences”, in accordance with the rules for admission to study in educational programs of higher education – bachelor’s programs, specialist’s programs, master’s programs at the Federal State Autonomous Educational Institution of Higher Education “Russian National Research Medical University named after N.I. Pirogov” of the Ministry of Health of the Russian Federation (FSAEIHE N.I. Pirogov RNIMU of the Ministry of Health of Russia) for the 2025/26 academic year.
II. Entry Test Program
Physical chemistry
1. Atomic structure
1.1 Particles in the atom and atomic radius
Candidates should be able to:
- understand that atoms are mostly empty space surrounding a very small, dense nucleus that contains protons and neutrons; electrons are found in shells in the empty space around the nucleus
- identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses
- understand the terms atomic and proton number; mass and nucleon number
- describe the distribution of mass and charge within an atom
- describe the behaviour of beams of protons, neutrons and electrons moving at the same velocity in an electric field
- determine the numbers of protons, neutrons and electrons present in both atoms and ions given atomic or proton number, mass or nucleon number and charge
- state and explain qualitatively the variations in atomic radius and ionic radius across a period and down a group
1.2 Isotopes
Candidates should be able to:
- define the term isotope in terms of numbers of protons and neutrons
- understand the notation xyА for isotopes, where x is the mass or nucleon number and y is the atomic or proton number
- state that and explain why isotopes of the same element have the same chemical properties
- state that and explain why isotopes of the same element have different physical properties, limited to mass and density
1.3 Electrons, energy levels and atomic orbitals
Candidates should be able to:
- understand the terms: shells, sub-shells and orbitals, principal quantum number, ground state, limited to electronic configuration
- describe the number of orbitals making up s, p and d sub-shells, and the number of electrons that can fill s, p and d sub-shells
- describe the order of increasing energy of the sub-shells within the first three shells and the 4s and 4p sub-shells
- describe the electronic configurations to include the number of electrons in each shell, sub-shell and orbital
- explain the electronic configurations in terms of energy of the electrons and inter-electron repulsion
- determine the electronic configuration of atoms and ions given the atomic or proton number and charge, using either of the following conventions: e.g. for Fe: 1s22s22p63s23p63d64s2 (full electronic configuration) or [Ar]3d64s2 (shorthand electronic configuration)
- understand and use the electrons in boxes notation
- describe and sketch the shapes of s and p orbitals
- describe a free radical as a species with one or more unpaired electrons
2 Atoms, molecules and stoichiometry
2.1 Relative masses of atoms and molecules
Candidates should be able to:
- define the unified atomic mass unit as one twelfth of the mass of a carbon-12 atom
- define relative atomic mass, Ar, relative isotopic mass, relative molecular mass, Mr, and relative formula mass in terms of the unified atomic mass unit
2.2 The mole and the Avogadro constant
Candidates should be able to:
- define and use the term mole in terms of the Avogadro constant
2.3 Formulae
Candidates should be able to:
- write formulae of ionic compounds from ionic charges and oxidation numbers
- write and construct equations (which should be balanced), including ionic equations
- define and use the terms empirical and molecular formula
- understand and use the terms anhydrous, hydrated and water of crystallisation
- calculate empirical and molecular formulae, using given data
2.4 Reacting masses and volumes (of solutions and gases)
Candidates should be able to:
- perform calculations including use of the mole concept, involving:
- reacting masses (from formulae and equations) including percentage yield calculations
- volumes of gases (e.g. in the burning of hydrocarbons)
- volumes and concentrations of solutions
- limiting reagent and excess reagent
- deduce stoichiometric relationships from calculations
3 Chemical bonding
3.1 Electronegativity and bonding
Candidates should be able to:
- define electronegativity as the power of an atom to attract electrons to itself
- explain the factors influencing the electronegativities of the elements in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells
- state and explain the trends in electronegativity across a period and down a group of the Periodic Table
- use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds
3.2 Ionic bonding
Candidates should be able to:
- define ionic bonding as the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions)
- describe ionic bonding
3.3 Metallic bonding
Candidates should be able to:
- define metallic bonding as the electrostatic attraction between positive metal ions and delocalised electrons
3.4 Covalent bonding and coordinate (dative covalent) bonding
Candidates should be able to:
- define covalent bonding as electrostatic attraction between the nuclei of two atoms and a shared pair of electrons
- describe covalent bonding in molecules
- understand that elements in period 3 can expand their octet
- describe coordinate (dative covalent) bonding
- describe covalent bonds in terms of orbital overlap giving σ and π bonds
- describe how the σ and π bonds form in molecules
- use the concept of hybridisation to describe sp, sp2 and sp3 orbitals
- define the terms:bond energy as the energy required to break one mole of a particular covalent bond in the gaseous state,
- bond length as the internuclear distance of two covalently bonded atoms
- use bond energy values and the concept of bond length to compare the reactivity of covalent molecules
3.5 Intermolecular forces, electronegativity and bond properties
Candidates should be able to:
- describe hydrogen bonding,
- limited to molecules containing N–H and O–H groups, including ammonia and water as simple examples
- use the concept of hydrogen bonding to explain the anomalous properties of H2O (ice and water): its relatively high melting and boiling points, its relatively high surface tension, the density of the solid ice compared with the liquid water
- use the concept of electronegativity to explain bond polarity and dipole moments of molecules
- describe van der Waals’ forces as the intermolecular forces between molecular entities other than those due to bond formation,
- and use the term van der Waals’ forces as a generic term to describe all intermolecular forces
- describe the types of van der Waals’ force
- state that, in general, ionic, covalent and metallic bonding are stronger than intermolecular forces
4 States of matter
4.1 Bonding and structure
Candidates should be able to:
- describe, in simple terms, the lattice structure of a crystalline solid which is:
- giant ionic, including sodium chloride and magnesium oxide
- simple molecular, including iodine, buckminsterfullerene C60 and ice
- giant molecular, including silicon(IV) oxide, graphite and diamond
- giant metallic, including copper
- describe, interpret and predict the effect of different types of structure and bonding on the physical properties of substances, including melting point, boiling point, electrical conductivity and solubility
- deduce the type of structure and bonding present in a substance from given information
5 Chemical energetics
5.1 Enthalpy change, ΔH
Candidates should be able to:
- understand that chemical reactions are accompanied by enthalpy changes and these changes can be exothermic (ΔH is negative) or endothermic (ΔH is positive)
- construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the reaction and of the activation energy
- define and use the terms:
- standard conditions (this syllabus assumes that these are 298 K and 101 kPa)
- enthalpy change with particular reference to: reaction, ΔHr , formation, ΔHf , combustion, ΔHc , neutralisation, ΔHneut
- understand that energy transfers occur during chemical reactions because of the breaking and making of chemical bonds
- use bond energies (ΔH positive, i.e. bond breaking) to calculate enthalpy change of reaction, ΔHr
- understand that some bond energies are exact and some bond energies are averages
5.2 Hess’s Law
Candidates should be able to:
- apply Hess’s Law to construct simple energy cycles
- carry out calculations using cycles and relevant energy terms, including:
- determining enthalpy changes that cannot be found by direct experiment
- use of bond energy data
6 Electrochemistry
6.1 Redox processes: electron transfer and changes in oxidation number (oxidation state)
Candidates should be able to:
- calculate oxidation numbers of elements in compounds and ions
- use changes in oxidation numbers to help balance chemical equations
- explain and use the terms redox, oxidation, reduction and disproportionation in terms of electron transfer and changes in oxidation number
- explain and use the terms oxidising agent and reducing agent
- indicate the magnitude of the oxidation number of an element
7 Equilibria
7.1 Chemical equilibria: reversible reactions, dynamic equilibrium
Candidates should be able to:
- understand what is meant by a reversible reaction
- understand what is meant by dynamic equilibrium in terms of the rate of forward and reverse reactions being equal and the concentration of reactants and products remaining constant
- understand the need for a closed system in order to establish dynamic equilibrium
- define Le Chatelier’s principle as: if a change is made to a system at dynamic equilibrium, the position of equilibrium moves to minimise this change
- use Le Chatelier’s principle to deduce qualitatively (from appropriate information) the effects of changes intemperature, concentration, pressure or presence of a catalyst on a system at equilibrium
- deduce expressions for equilibrium constants in terms of concentrations, Kc
- use the terms mole fraction and partial pressure
- deduce expressions for equilibrium constants in terms of partial pressures, Kp (use of the relationship between Kp and Kc is not required)
- use the Kc and Kp expressions to carry out calculations
- calculate the quantities present at equilibrium, given appropriate data
- state whether changes in temperature, concentration or pressure or the presence of a catalyst affect the value of the equilibrium constant for a reaction
7.2 Arrhenius and Brønsted–Lowry theory of acids and bases
Candidates should be able to:
- state the names and formulae of the common acids
- state the names and formulae of the common alkalis
- describe the Brønsted–Lowry theory of acids and bases
- describe strong acids and strong bases as fully dissociated in aqueous solution and weak acids and weak bases as partially dissociated in aqueous solution
- appreciate that water has pH of 7, acid solutions pH of below 7 and alkaline solutions pH of above 7
- explain qualitatively the differences in behaviour between strong and weak acids including the reaction with a reactive metal and difference in pH values by use of a pH meter, universal indicator or conductivity
- understand that neutralisation reactions occur when H+(aq) and OH–(aq) form H2O(l)
- understand that salts are formed in neutralisation reactions
8 Reaction kinetics
8.1 Rate of reaction
Candidates should be able to:
- explain and use the term rate of reaction, frequency of collisions, effective collisions and non-effective collisions
- explain qualitatively, in terms of frequency of effective collisions, the effect of concentration and pressure changes on the rate of a reaction
- use experimental data to calculate the rate of a reaction
8.2 Effect of temperature on reaction rates and the concept of activation energy
Candidates should be able to:
- define activation energy, EA, as the minimum energy required for a collision to be effective
8.3 Homogeneous and heterogeneous catalysts
Candidates should be able to:
1) explain and use the terms catalyst and catalysis
- explain that, in the presence of a catalyst, a reaction has a different mechanism, i.e. one of lower activation energy
- construct and interpret a reaction pathway diagram, for a reaction in the presence and absence of an effective catalyst
Inorganic chemistry
9 The Periodic Table: chemical periodicity
9.1 Periodicity of physical properties of the elements in Period 3
Candidates should be able to:
- describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements
- explain the variation in melting point and electrical conductivity in terms of the structure and bonding of the еlements
9.2 Periodicity of chemical properties of the elements in Period 3
- describe, and write equations for, the reactions of the elements with oxygen
- state and explain the variation in the oxidation number of the oxides and chlorides in terms of their outer shell (valence shell) еlectrons
- describe, and write equations for, the reactions, if any, of the oxides with water including the likely pHs of the solutions obtained
- describe, explain, and write equations for, the acid / base behaviour of the oxides and the hydroxides including, where relevant, amphoteric behavior in reactions with acids and bases
- describe, explain, and write equations for, the reactions of the chlorides with water including the likely pHs of the solutions obtained
- explain the variations and trends in terms of bonding and electronegativity
- suggest the types of chemical bonding present in the chlorides and oxides from observations of their chemical and physical properties
9.3 Chemical periodicity of other elements
Candidates should be able to:
- predict the characteristic properties of an element in a given group by using knowledge of chemical periodicity
- deduce the nature, possible position in the Periodic Table and identity of unknown elements from given information about physical and chemical properties
10 Group 2
10.1 Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their сompounds
Candidates should be able to:
- describe, and write equations for, the reactions of the elements with oxygen, water and dilute hydrochloric and sulfuric acids
- describe, and write equations for, the reactions of the oxides, hydroxides and carbonates with water and dilute hydrochloric and sulfuric acids
- describe, and write equations for, the thermal decomposition of the nitrates and carbonates, to include the trend in thermal stabilities
11 Group (7) 17
11.1 Physical properties of the Group 17 elements
Candidates should be able to:
- describe the colours and the trend in volatility of chlorine, bromine and iodine
- describe and explain the trend in the bond strength of the halogen molecules
- interpret the volatility of the elements in terms of instantaneous dipole–induced dipole forces
11.2 The chemical properties of the halogen elements and the hydrogen halides
Candidates should be able to:
- describe the relative reactivity of the elements as oxidising agents
- describe the reactions of the elements with hydrogen and explain their relative reactivity in these reactions
- describe the relative thermal stabilities of the hydrogen halides and explain these in terms of bond strengths
- describe the relative reactivity of halide ions as reducing agents
11.3 The reactions of chlorine
Candidates should be able to:
- describe and interpret, in terms of changes in oxidation number, the reaction of chlorine with cold and with hot aqueous sodium hydroxide and recognise these as disproportionation reactions
12 Nitrogen and sulfur
12.1 Nitrogen and sulfur
Candidates should be able to:
- explain the lack of reactivity of nitrogen, with reference to triple bond strength and lack of polarity
- describe and explain:
- the basicity of ammonia, using the Brønsted–Lowry theory
- the structure of the ammonium ion and its formation by an acid–base reaction
- the displacement of ammonia from ammonium salts by an acid–base reaction
- state and explain the natural and man-made occurrences of oxides of nitrogen and their catalytic removal from the exhaust gases of internal combustion engines
- describe the role of NO and NO2 in the formation of acid rain both directly and in their catalytic role in the oxidation of atmospheric sulfur dioxide
Organic chemistry
13 An introduction to organic chemistry
13.1 Formulae, functional groups and the naming of organic compounds
Candidates should be able to:
- define the term hydrocarbon as a compound made up of C and H atoms only
- understand that alkanes are simple hydrocarbons with no functional group
- understand that the compounds contain a functional group which dictates their physical and chemical properties
- interpret and use the general, structural, displayed and skeletal formulae of the classes of compound
- understand and use systematic nomenclature of simple aliphatic organic molecules with functional groups
- deduce the molecular and/or empirical formula of a compound, given its structural, displayed or skeletal formula
13.2 Characteristic organic reactions
Candidates should be able to:
- interpret and use the following terminology associated with types of organic compounds and reactions:
- homologous series
- saturated and unsaturated
- homolytic and heterolytic fission
- free radical, initiation, propagation, termination
- nucleophile, electrophile, nucleophilic, electrophilic
- addition, substitution, elimination, hydrolysis, condensation
- oxidation and reduction
- understand and use the following terminology associated with types of organic mechanisms:
- free-radical substitution
- electrophilic addition
- nucleophilic substitution
- nucleophilic addition
13.3 Shapes of organic molecules; σ and π bonds
Candidates should be able to:
- describe organic molecules as either straight-chained, branched or cyclic
- describe and explain the shape of, and bond angles in, molecules containing sp, sp2 and sp3 hybridised atoms
- describe the arrangement of σ and π bonds in molecules containing sp, sp2 and sp3 hybridised atoms
- understand and use the term planar when describing the arrangement of atoms in organic molecules, for example ethane
13.4 Isomerism: structural and stereoisomerism
Candidates should be able to:
- describe structural isomerism and its division into chain, positional and functional group isomerism
- describe stereoisomerism and its division into geometrical (cis/trans) and optical isomerism
- describe geometrical (cis/trans) isomerism in alkenes, and explain its origin in terms of restricted rotation due to the presence of π bonds
- explain what is meant by a chiral centre and that such a centre gives rise to two optical isomers (enantiomers)
- identify chiral centres and geometrical (cis/trans) isomerism in a molecule of given structural formula including cyclic compounds
- deduce the possible isomers for an organic molecule of known molecular formula
14 Hydrocarbons
14.1 Alkanes
Candidates should be able to:
- recall the reactions (reagents and conditions) by which alkanes can be produced:
- addition of hydrogen to an alkene in a hydrogenation reaction, H2(g) and Pt/Ni catalyst and heat
- cracking of a longer chain alkane, heat with Al2O3
- describe:
- the complete and incomplete combustion of alkanes
- the free-radical substitution of alkanes by Cl2 or Br2 in the presence of ultraviolet light, as exemplified by the reactions of ethane
- describe the mechanism of free-radical substitution with reference to the initiation, propagation and termination steps
- suggest how cracking can be used to obtain more useful alkanes and alkenes of lower Mr from heavier crude oil fractions
- understand the general unreactivity of alkanes, including towards polar reagents in terms of the strength of the C–H bonds and their relative lack of polarity
14.2 Alkenes
Candidates should be able to:
- recall the reactions (including reagents and conditions) by which alkenes can be produced:
- elimination of HX from a halogenoalkane by ethanolic NaOH and heat
- dehydration of an alcohol, by using a heated catalyst (e.g. Al2O3) or a concentrated acid
- cracking of a longer chain alkane
- describe the following reactions of alkenes:
- the electrophilic addition of :
- hydrogen in a hydrogenation reaction, H2(g) and Pt/Ni catalyst and heat
- steam, H2O(g) and H3PO4 catalyst
- a hydrogen halide, HX(g) at room temperature
- a halogen, X2
- the electrophilic addition of :
-
- the oxidation by cold dilute acidified KMnO4 to form the diol
- the oxidation by hot concentrated acidified KMnO4 leading to the rupture of the carbon–carbon double bond and the identities of the subsequent products to determine the position of alkene linkages in larger molecules
- addition polymerisation exemplified by the reactions of ethene and propene
- describe the use of aqueous bromine to show the presence of a C=C bond
- describe the mechanism of electrophilic addition in alkenes, using bromine / ethene and hydrogen bromide / propene as examples
- describe and explain the inductive effects of alkyl groups on the stability of primary, secondary and tertiary cations formed during electrophilic addition (this should be used to explain Markovnikov addition)
14.3 Cycloalkanes
Candidates should be able to:
- describe and explain structure, homologous series, nomenclature, isomerism.
- describe and explain chemical properties: most common addition reactions: halogenation, addition of alkyl halides, hyration, hydratation, nitration.
14.4 Aromatic hydrocarbons. Arenes
Candidates should be able to:
- describe and explain chemical and electronic structures of benzene. Benzene- cyclic conjugated system. Conjugation energy. Homologous series of benene, nomenclature, isomerism.
- describe and explain chemical properties of benzene: reactions of electrophilic substitution (nitration, sulfonation, halogenation, alkylation – with halogenated alkanes, reactions with alkenes; acylation).
- describe and explain electrophilic substitution mechanism.
- describe and explain addition reactions (addition of hydrogen, halogens).
- describe and explain chemical properties of benzene homologues. Mutual influence of atoms in cyclic hydrocarbons. Orientation in benzene rings.
- describe and explain oxidation reaction.
15 Halogen compounds
15.1 Halogenoalkanes
Candidates should be able to:
- recall the reactions (reagents and conditions) by which halogenoalkanes can be produced:
- the free-radical substitution of alkanes by Cl2 or Br2 in the presence of ultraviolet light, as exemplified by the reactions of ethane
- electrophilic addition of an alkene with a halogen, X2, or hydrogen halide, HX(g), at room temperature
- substitution of an alcohol, e.g. by reaction with HX or KBr with H2SO4 or H3PO4; or with PCl3 and heat; or with PCl5; or with SOCl2
- classify halogenoalkanes into primary, secondary and tertiary
- describe the following nucleophilic substitution reactions:
- the reaction with NaOH(aq) and heat to produce an alcohol
- the reaction with KCN in ethanol and heat to produce a nitrile
- the reaction with NH3 in ethanol heated under pressure to produce an amine
- the reaction with aqueous silver nitrate in ethanol as a method of identifying the halogen present as exemplified by bromoethane
- describe the elimination reaction with NaOH in ethanol and heat to produce an alkene as exemplified by bromoethane
- describe the SN1 and SN2 mechanisms of nucleophilic substitution in halogenoalkanes including the inductive effects of alkyl groups
- recall that primary halogenoalkanes tend to react via the SN2 mechanism; tertiary halogenoalkanes via the SN1 mechanism; and secondary halogenoalkanes by a mixture of the two, depending on structure
- describe and explain the different reactivities of halogenoalkanes (with particular reference to the relative strengths of the C–X bonds as exemplified by the reactions of halogenoalkanes with aqueous silver nitrates)
16 Hydroxy compounds
16.1 Alcohols
Candidates should be able to:
- recall the reactions (reagents and conditions) by which alcohols can be produced:
- electrophilic addition of steam to an alkene, H2O(g) and H3PO4 catalyst
- reaction of alkenes with cold dilute acidified potassium manganate(VII) to form a diol
- substitution of a halogenoalkane using NaOH(aq) and heat
- reduction of an aldehyde or ketone using NaBH4 or LiAlH4
- reduction of a carboxylic acid using LiAlH4
- hydrolysis of an ester using dilute acid or dilute alkali and heat
- describe:
- the reaction with oxygen (combustion)
- substitution to halogenoalkanes, e.g. by reaction with HX or KBr with H2SO4 or H3PO4; or with PCl3 and heat; or with PCl5; or with SOCl2
- the reaction with Na(s)
- oxidation with acidified K2Cr2O7 or acidified KMnO4 to:
- carbonyl compounds by distillation
- carboxylic acids by refluxing (primary alcohols give aldehydes which can be further oxidised to carboxylic acids, secondary alcohols give ketones, tertiary alcohols cannot be oxidised)
- dehydration to an alkene, by using a heated catalyst, e.g Al2O3or a concentrated acid
- formation of esters by reaction with carboxylic acids and concentrated H2SO4 or H3PO4 as catalyst as exemplified by ethanol
- classify alcohols as primary, secondary and tertiary alcohols, to include examples with more than one alcohol group
- state characteristic distinguishing reactions, e.g. mild oxidation with acidified K2Cr2O7, colour change from orange to green
- explain the acidity of alcohols compared with water
16.2 Phenols
Candidates should be able to:
- describe and explain structure of phenols, nomenclature and isomerism
- describe and explain chemical properties: acidic properties, reactions of electrophilic substitution in benzene rings (nitration, sulfonation, reaction with bromine water), reduction reactions.
17 Carbonyl compounds
17.1 Aldehydes and ketones
Candidates should be able to:
- recall the reactions (reagents and conditions) by which aldehydes and ketones can be produced:
- the oxidation of primary alcohols using acidified K2Cr2O7 or acidified KMnO4 and distillation to produce aldehydes
- the oxidation of secondary alcohols using acidified K2Cr2O7 or acidified KMnO4 and distillation to produce ketones
- describe:
- the reduction of aldehydes and ketones, using NaBH4 or LiAlH4to produce alcohols
- the reaction of aldehydes and ketones with HCN, KCN as catalyst, and heat to produce hydroxynitriles
- deduce the nature (aldehyde or ketone) of an unknown carbonyl compound from the results of simple tests (Fehling’s and Tollens’ reagents; ease of oxidation)
18 Carboxylic acids and derivatives
18.1 Carboxylic acids
Candidates should be able to:
- recall the reactions by which carboxylic acids can be produced:
- oxidation of primary alcohols and aldehydes with acidified K2Cr2O7 or acidified KMnO4 and refluxing
- hydrolysis of nitriles with dilute acid or dilute alkali followed by acidification
- hydrolysis of esters with dilute acid or dilute alkali and heat followed by acidification
- describe:
- the redox reaction with reactive metals to produce a salt and H2(g)
- the neutralisation reaction with alkalis to produce a salt and H2O(l )
- the acid–base reaction with carbonates to produce a salt and H2O(l) and CO2(g)
- esterification with alcohols with concentrated H2SO4 as catalyst
- reduction by LiAlH4 to form a primary alcohol
18.2 Esters
Candidates should be able to:
- recall the reaction (reagents and conditions) by which esters can be produced:
- the condensation reaction between an alcohol and a carboxylic acid with concentrated H2SO4 as catalyst
- describe the hydrolysis of esters by dilute acid and by dilute alkali and heat
19 Nitrogen compounds
19.1 Primary amines
Candidates should be able to:
- recall the reactions by which amines can be produced:
- reaction of a halogenoalkane with NH3 in ethanol heated under pressure
19.2 Nitriles and hydroxynitriles
Candidates should be able to:
- recall the reactions by which nitriles can be produced:
- reaction of a halogenoalkane with KCN in ethanol and heat
- recall the reactions by which hydroxynitriles can be produced:
- the reaction of aldehydes and ketones with HCN, KCN as catalyst, and heat
- describe the hydrolysis of nitriles with dilute acid or dilute alkali followed by acidification to produce a carboxylic acid
20 Polymerisation
20.1 Addition polymerisation
Candidates should be able to:
- describe addition polymerisation as exemplified by poly(ethene) and poly(chloroethene), PVC
- deduce the repeat unit of an addition polymer obtained from a given monomer
- identify the monomer(s) present in a given section of an addition polymer molecule
- recognise the difficulty of the disposal of poly(alkene)s, i.e. non-biodegradability and harmful combustion рroducts
21 Organic synthesis
21.1 Organic synthesis
Candidates should be able to:
- for an organic molecule containing several functional groups:
- identify organic functional groups using the reactions in the syllabus
- predict properties and reactions
- devise multi-step synthetic routes for preparing organic molecules using the reactions in the syllabus
- analyse a given synthetic route in terms of type of reaction and reagents used for each step of it, and possible by-products