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Olga V. Kovalchukova, Nasrin Namichemazi & Rusul Alabada Lectures in CHEMISTRY For the 1-st year students of the «Dentistry» speciality of the Medical Faculty of the Peoples’ Friendship University of Russia Ковальчукова О.В., Насрин Намичемази, Русул Алабада ХИМИЯ (конспект лекций) Для студентов 1 курса медицинского факультета специальности «Стоматология» (на английском языке) Москва Издательство Российского университета дружбы народов, 2015 Утверждено РИС Ученого совета Российского университета дружбы народов Ковальчукова О.В., Насрин Намичемази, Русул Алабада. Химия. Конспект лекций для студентов I курса медицинского факультета специальности «Стоматология». М.: Изд-во РУДН, 2015. – 164 с. Настоящее учебное пособие представляет собой конспект лекций, предназначенный для студентов медицинского факультета специальности «Стоматология». Составлено в соответствии с Федеральным государственным образовательным стандартом и программой курса «Химия». Предназначено для работы студентов I курса медицинского факультета специальности «Стоматология» при подготовке к лабораторным занятиям и к экзамену по курсу «Химия». Подготовлено на кафедре общей химии. The manual represents the lecture book of the Chemistry discipline delivering for the students of the «Dentistry» speciality of the Medical Faculty of Peoples’ Friendship University of Russia. It is prepared in accordance with the Federal state educational standard and the Program in Chemistry. Intended for the 1-st year students of the «Dentistry» speciality of the Medical Faculty of the Peoples’ Friendship University of Russia. Prepared at the department of General Chemistry. © Ковальчукова О.В., Насрин Намичемази, Русул Алабада © Издательство Российского университета дружбы народов, 2015 2 SUMMARY Page Part 1. General Chemistry Introduction………………………………………………………. General notions of Chemistry…………………………………… General Laws of Chemistry……………………………………... Classes of Inorganic Compounds……………………………….. Concepts of Chemical Thermodynamics………………………… Chemical Kinetics….……………………………………………. Structure of Atoms……………………………………………… Periodic Law and Periodic System of Chemical Elements………. Chemical bonds…………………………………………………… Adsorption…………………………………………..………….…. Concepts of Catalysis………………………………………….….. Solutions………………………………………..………………… Physico-Chemical Processes in solutions…………..……………. Dissociation of Strong Electrolytes…...………………………… Dissociation of weak electrolytes…………………………………. Electrolytic Dissociation of Water……………………………… Equilibria in Solutions with Precipitates..……………………... Buffer solutions…………………………………………………. Direction of reactions of ionic exchange………………………… Hydrolysis of salts………………………………………………… Colloidal solutions……………………………………………….. Oxidation-reduction reaction (redox-reactions)………………… Complex compounds…………………………………………….. 3 5 7 9 15 18 21 28 30 40 43 45 46 49 50 51 52 54 57 59 64 70 76 Part 2. Inorganic Chemistry s-Elements (alkaline and alkaline earth metals)………………… Elements of IIIA and IVA groups (p-elements)…………………. Elements of VA and VIA groups (p-elements)…………………. Elements of the VIIA group (halogens)………………………….. Transitional elements (d-elements)………………………………. 85 87 89 90 92 3 Electrochemical corrosion of metals…………………………….. 93 Part 3. Organic Chemistry The theory of chemical structure of organic compounds………... Alkanes……………………………………………………………. Cycloalkanes……………………………………………………… Alkenes……………………………………………………………. Alkadienes………………………………………………………… Alkynes……………………………………………………………. Aromatic hydrocarbons…………………………………………... Organic compounds with functional groups…………………….. Amines……………………………………………………………. Alcohols…………………………………………………………… Phenols…………………………………………………………… Aldehydes and ketones……………………………………………. Carboxylic acids…………………………………………………... Aminoacids and proteins………………………………………… Esters, fats and oils………………………………………………. Carbohydrates……………………………………………………. Sulphur in organic compounds…………………………………. Biologically important heterocyclic compounds………………… List of citated literature…………………………………………. 4 97 106 108 108 110 112 113 118 119 122 128 130 135 142 145 148 156 157 164 121 121 125 128 131 136 142 146 149 157 157 PART 1 GENERAL CHEMISTRY 5 6 BASIC CONCEPTS OF CHEMISTRY Objects of studying of chemistry are substances and their smallest particles - molecules and atoms. Molecule is the least particle of substance possessing its chemical properties. Atom is the least particle of a chemical element possessing its chemical properties. Atoms are components of molecules. Chemical element is the kind of atoms characterized by certain set of properties. Mole is a unit of measurement of quantity of substances, containing such amount of molecules, atoms or other structural units, as 12 g of an isotope 12С. The number of structural units containing in 1 mole of substances, is known as Avogadro’s number (NA): NA = 6,021023 mol-1. The mass of 1 mole of a substance expressed in grams, is called a molar mass of a substance (M, g x mol-1). Quantity of a substance: = m / M = N / NA (mole). GENERAL LAWS OF CHEMISTRY 1. Incorporated gas law: PV P0 V0 T T0 . 2. Avogadro’s Law: equal volumes of various gases under identical conditions (temperature and pressure) contain the identical number of molecules. Consequence 1. Masses of two identical volumes of various gases under identical conditions concern as their molar masses: 7 m1 M 1 D m2 M2 (relative density of a gas).. Consequence 2. Under normal conditions (Р0 = 101325 Pa and Т0 = 273 K) one mole of any gas occupies the volume of 22,4 l (VМ - molar volume). 3. Mendeleyev-Klayperon’s equation: PV m RT M , where R - a universal gas constant (R = 8,314 J x mole-1 x K). 4. Partial pressure of gas is the part of a total pressure of a gas mixture which is necessary on a share of the given gas. According to Dalton’s law, partial pressure of gas in a mixture equals the pressure of gas as if it occupied the total volume under the same temperature. Рtot. = Р1 + Р2 + … + Рn. 5. The law of mass conservation. Sum of masses of the substances entering a chemical reaction, equal total mass of the products. 6. The law of equivalents. Substances react with each other in the quantities proportional to their equivalents: n1 = n2 (n - number of equivalents). The equivalent is such a quantity of substance which reacts with 1 mole of hydrogen atoms or replaces them in chemical reactions (the quantity of substance corresponding to a unit valency). Equivalent mass (МE) is the mass of one equivalent of the substance, expressed in grams: МE = f x M (where f is the factor of equivalence). 8 Calculation of the factor of equivalence of different classes of inorganic compounds: For simple substances and elements in chemical compounds f = 1 / V (where V is the valency of an element). For acids and the bases f = 1 / m (where m is the basidity of an acid or acidity of a base). For oxides and salts f = 1 / n x V (where n is the number of metallic atoms in the compound, and V is the valency of the metal). Number of equivalents: n = m / МE (for any substance); n = V / VE (for gaseous substances), VE is the equivalent volume of the gas (the volume occupied by one equivalent of a gas under normal conditions). For example, under normal conditions the equivalent volume of hydrogen (МE = 1 g x mole-1) equals 11,2 liters, and equivalent volume of oxygen (МE = 8 g x mole-1) equals 5,6 litres. GENERAL CLASSES OF INORGANIC COMPOUNDS OXIDES Oxides are binary compounds containing oxygen atoms in the –2 oxidation state. 9 OXIDES Salt-forming Not salt-forming (indifferent) Basic Amphoteric Acidic Metallic oxides in low oxidation states (+1, +2) Some metallic oxides in the oxidation states +2, +3 and +4 Non-metallic oxides in oxidation states +3 and above, as far as metallic oxides in oxidation states +4 and above Non-metallic oxides in oxidation states +1 and+2 Na2O, Cu2O, FeO, MgO BeO, ZnO, PbO, SnO, Al2O3, Cr2O3, MnO2 В2О3, СО2, P2O5, PbO2, CrO3 СО, SiO, N2O, NO Increase in oxidation state of a non-metal changes properties of its oxide from indifferent to acidic: N2O NO N2O3 NO2 N2O5 Indifferent oxides Acidic oxides Increase in oxidation state of a metal changes properties of its oxide from basic via amphoteric to acidic: MnO Mn2O3 MnO2 MnO3 Mn2O7 Basic oxides Amphoteric Acidic oxides oxide Chemical properties of oxides: 10 1. Basic and acidic oxides are dissolved in water (react with it) in case if soluble bases and acids are formed:: Na2O + H2O = 2NaOH Only basic oxides of alkaline and CaO + H2O = Ca(OH)2 alkaline-earth metals FeO + H2O SO3 + H2O = H2SO4 P2O5 + H2O (хол.) = 2НРО3 P2O5 + 3H2O (гор.) = 2Н3РО4 2NO2 + H2O = HNO2 + HNO3 4NO2 + O2 + 2H2O = 4HNO3 СrO3 + H2O = H2CrO4 SiO2 + H2O Only acidic oxides which form soluble acids 2. Acidic oxides react with basic oxides: Na2O + СО2 = Na2СО3 If the acidic oxide is in the gaseous phase, the reaction takes CaO + SO2 = CaSO3 place under room temperature melting If the acidic oxide is solid, the Na2O + SiO2 Na2SiO3 reaction takes place undervhigh melting temperatures (melting) FeO + PbO2 FePbO3 3. Amphoteric oxides can react with both acidic and basic oxides: ZnO + Na2O Na2ZnO2 ZnO + CO2 = ZnCO3 melting melting 2KAlO2 melting Al2O3+3SiO2 Al2(SiO3)3 Al2O3 + K2O 11 BASES Bases are substances, which form hydroxyl-anions (OH-) at dissociation. Number of OH-groups of the base defines its acidity. Chemical properties of bases: Alkalis: 1) react with acidic oxides and acids: 2КОН + SO3 = K2SO4 + H2O Ca(OH)2 + 2HCl = CaCl2 + 2H2O Reaction of neutralization 2) react with salts (in case if a precipitate is formed): 2NaOH + FeCl2 = Fe(OH)2 + 2NaCl Ba(OH)2 + Na2SO4 = BaSO4 + 2NaOH NaOH + BaCl2 3) react with amphoteric oxides and amphoteric bases: 2NaOH + BeO 3NaOH + Cr(OH)3 melting Na2BeO2 + H2O solution Na3[Cr(OH)6] 4) dissociate in aqueous solutions and change the colour of acidbase indicators: NaOH = Na+ + OH– (methylorange turns yellow, lithmuth turns blue, and phenolphtalein turns pink) 12 Insoluble bases: 1) react with acidic oxides and acids (reaction of neutralization): Fe(OH)3 + 3HNO3 = Fe(NO3)3 + 3H2O t 2) decompose at heating: Cu(OH)2 CuO + H2O Amphoteric bases (bases which correspond to amphoteric oxides): 1) react with acidic oxides and acids (reaction of neutralization): Zn(OH)2 + CO2 = ZnCO3 + H2O Al(OH)3 + 3HCl = AlCl3 + 3H2O 2) react with alkalis: Sn(OH)2 + 2NaOH = Na2[Sn(OH)4] 3) decompose at heating: t 2Al(OH)3 Al2O3 + 3H2O ACIDS Acids are substances, which form hydrogen kations (H+) at dissociation. Basicity of an acid (number of H-atoms) defines the possibility of full or incomplete neutralisation of an acid in reactions with bases: HCl + NaOH = NaCl + H2O (for monobasic acids only one reaction of neutralisation is possible). For the multibasic acids full and incomplete neutralisation is possible: H2SО4 + 2КОН = К2SО4 + 2Н2О (full neutralisation) H2SО4 + КОН = КНSО4 + Н2О (incomplete neutralization) 13 Chemical properties of acids: 1. 2. 3. Change in colour of indicators (methylorange and lithmus turn red). Interaction with active metals: Са + 2HCl = CaCl2 + H2 Interaction with basic and amphoteric oxides: СuO + 2HNO3 = Cu(NO3)2 + H2O Al2O3 + 3H2SO4 = Al2(SO4)3 + 3H2O 4. Interaction with bases (reaction of neutralization): Cu(OH)2 + 2HCl = CuCl2 + 2H2O 5. Interaction with salts (reactions of ionic exchange): H2SO4 + BaCl2 = 2HCl + BaSO4 (precipitate is formed) 2HNO3 + СаСО3 = Ca(NO3)2 + H2O + CO2 (gas is evolved) HCl + KNO2 = KCl + HNO2 (weak acid is formed) SALTS Salts are substances, which form metallic kations or ammonium (NH4+) and acidic anions at dissociation. Salts can be considered as products of replacement of hydrogen atoms in an acid molecule by metals or hydroxyl-groups in a base by an acidic anion. Chemical properties of salts: 1. 2. 3. 14 Interaction with metals (a more active metal replaces a less active one): FeCl2 + Zn = Fe + ZnCl2. Interaction with non-metals (a more active non-metal replaces a less active one): 2NaBr + Сl2 = Br2 + 2NaCl. Interaction with alkalis (reactions of ionic exchange): MgCl2 + 2NaOH (exsess) = Mg(OH)2 + 2NaCl. MgCl2 + NaOH (lack) = MgOHCl + NaCl. 4. Interaction with acids (reactions of ionic exchange): СаCl2 + H2SO4 (lack) = CaSO4 + 2HCl. CaSO4 + H2SO4 (exsess) = Ca(HSO4)2 5. Interaction between two salts (reactions of ionic exchange): CuCl2 + 2AgNO3 = 2AgCl + Cu(NO3)2. CONCEPTS OF CHEMICAL THERMODYNAMICS The study of chemical processes should be approached through a series of successive approximations. The first stage is expedient to consider only the initial and final states of the interacting bodies without considering the way in which the process takes place, and the development of the process in time. This approach is called the thermodynamics. For convenience, the objects of study should be isolated. Such a collection of bodies extracted from the space, is called the system . If no mass and heat transfer exists between the system and the surrounding environment, the system is called isolated. If this condition is not met, then the system is called open. If the system is only possible for the heat transfer , it is called closed. A state of any system is characterized by certain thermodynamic parameters , which include temperature (T ) , pressure ( P) , volume (V), chemical composition . The change of at least one of the parameters leads to a change in the system state. State of the system can be represented in the form of so-called state equation: ( P , V, T) = 0 For majority of systems, the state equations are empirical , that is, the experimentally obtained equations describing the behavior of matter in a certain range of pressure and temperature. 15 For the thermodynamic description of the system, the state functions are used. These are equations that can be uniquely identified by the parameters P, V and T. The values of these functions are independent of the nature of the process, resulting in a system of this state . The functions of the state are: 1) the internal energy of the system (U); 2) the enthalpy ( heat content ) of ( H) ; 3 ) the entropy (a measure of disorder) of the system (S); 4) the Gibbs’ free energy (G); 5) the Helmholtz’s free energy (F). Chemical reactions are accompanied by the release or absorption of energy usually in the form of heat. Th reactions in which heat is released are called exothermic reactions and extending the absorption of heat endothermic. Since the heat generation decreases the enthalpy of the system, then Q = - H, where Q is the heat of the reaction, and H is the change in enthalpy of the system. Thus, the condition of the exothermic reaction is Q> 0 or H < 0, and the condition of the endothermic reaction - Q < 0 or H> 0 . The equation of a chemical reaction involving the magnitude of the thermal effect (enthalpy) is called a thermochemical equation: 2H2 ( g ) + O2 ( g ) = 2H2O ( g) + 571.6 kJ or 2H2 ( g ) + O2 ( g ) = 2H2O ( g) ; H = 571.6 kJ Heat of formation of compounds is the amount of heat released during the formation of 1 mole of the compound from elements in their most stable modification . Thus, the heat of formation of water Hf (H2O ) = 571.6 / 2 = 285.8 kJ / mol The heat of formation of the substance, measured in standard conditions (T = 298 K, P = 101325 Pa), is called the standard heat of formation , and is denoted H0. The standard heat of formation of simple substance in its most stable modification shall be equal to zero. Calculation of the heat of the reaction from the heats of formation of the participating substances, is produced by the Hess’ law , the heat of the chemical reaction depends on the starting and final products state and is not dependent on any stage after the reaction proceeds . The thermal effect 16 of the process is the sum of the thermal effects of the individual stages of the process. For example, coal combustion may take place in a single step: C (s) + O2 (g) = CO2 (g); H1 = 395.4 kJ or through the intermediate formation of carbon oxide (II): a) C (s) + 1 / 2O2 (g) = CO (g); H2 = 110.7 kJ b) CO (g) + 1 / 2O2 (g) = CO2 (g); H3 = 284.7 kJ . The total heat released by the reaction in both cases is the same: H1 = H2 + H3. By the corollary of the Hess’ law, the heat of the reaction is equal to the difference between the sum of the standard heats of formation of the final products and raw materials. For example, the reaction MgO (s) + CO2 (g) = MgCO3 (s) H0 reaction = H0 (MgCO3) - [H0 (MgO) + H0 (CO2)] = = 115.6 - (- 602.0 - 395.4 ) = 1113.0 kJ . The standard heats of formation of substances are tabulated data. There are two driving forces for the nature of the processes occurring in the nature. These are the desire to move to the lower energy state (H < 0) and the desire to move into a state of the greatest disorder (entropy) (S> 0). Since chemical reactions usually proceed with the change of the energy of the system and its entropy, then the reaction’s direction that where the total driving force of the reaction decreases. Under the isobaric- isothermal conditions (pressure and temperature do not change) the total driving force of the reaction is called the Gibbs’ energy: G = H - TS The negative value of the Gibbs’ energy change (G < 0) is the condition for the spontaneous reaction. The temperature at which G = 0, is called the reaction initiation temperature. In this case TG = 0 = H / S. 17 Changes in the Gibbs’ energy and entropy in chemical reactions are similar to the change in the enthalpy (heat effect) and determined in accordance with the result of the Hess’ law: H0 = (H0products - H0 . reactants) G0 = (G0products - G0 . reactants) S0 = (S0products - S0. reactants) . CHEMICAL KINETICS One of the basic concepts of chemical kinetics is the concept of a rate of a chemical reaction. Rate of a chemical reaction is denoted as number of elementary acts of a reaction which results transformation of reactants into reaction products, in a unit time in a unit volume. In practice, rates of reactions can be measured as a change in concentrations of substances participating in it for a certain time interval: v c Out of two chemical reactions, that one is of the greatest rate, in which under identical time more quantity of a substance is formed. Law of mass action. Collision of molecules should be a necessary condition for realisation of chemical interaction between molecules. Collision occurs the more often, than more molecules contains in the given volume, i.e. rate of a chemical reaction depends on concentrations of reacting substances. aA bB mM nN v k A B where k is a rate constant of the chemical reaction, numerically it equals rate of a reaction at unit concentrations of reacting substances. , - simple numbers, usually not more than 3. For simple reactions they correspond to stoitiometric coefficients of the reaction. 18 The rate of the reaction does not depend on concentrations of firm substances, but only on their surface area. CaO + CO2 CaCO3 v k CO2 v k 2 The equation of a reaction frequently does not reflect its mechanism. For example, the reaction 2HI + H2O2 2H2O + I2 really proceeds in two stages: 1. HI + H2O2 HOI + H2O (slowly) 2. HOI + HI I2 + H2O (quickly) Kinetics of the overall reaction it is described by the first (slow) stage. Expression of speed of this reaction registers as v k HI H 2 O2 instead of , v k HI H 2 O2 2 Temperature dependence of rates of chemical reactions. Rate of a chemical reaction depends on number of effective collisions. Effective collision occurs only between active molecules. Increase in temperature increases number of active molecules, providing them with necessary activation energy, and the rate of the reaction increases. Activation energy is that additional energy which it is necessary to transfer to system to start chemical reaction. Vant Hoff’s rule. At increase in temperature on 10 speed of reaction increases in 2 - 4 times. T2 T1 v 2 v1 10 , v2 v1 where is the rate of a reaction at temperature T2, is the rate of a reaction at temperature T1, is the temperature coefficient of the reaction which defines change of the rate of the reaction at temperature change on 10. 19 v As the reactional time. c T2 T1 1 2 10 , where so is Chemical equilibrium If a chemical reaction can proceed only in one direction it is called as irreversible. The reactions proceeding simultaneously in two directions, are reversible. Eventually rate of a direct reaction (v ) decreases, and rate of a back reaction (v ) increases until they become equal. So, a chemical equilibrium is established in the system. The condition of a chemical equilibrium: v = v . In the equilibrium state, reversible reactions are described by an equilibrium constant K: aA + bB K c C a A C C C C d D b B cC + dD where CA, CB, CC, CD are concentrations of gaseous or dissolved substances. Chemical equilibrium is a dynamic one, so it can be shifted according to le Chateleu’s principle (principle of counteraction): if an equilibrium system is affected by any factor (change in concentrations, pressure or temperature), the equilibrium will be shifted in the direction which weakens the external influence. The increase in temperature shifts the equilibrium towards an endothermic reaction (the system absorbs heat and increases its internal energy, H>0), and decrease in temperature shifts the equilibrium towards 20 an exothermic reaction (the system evolves heat and decreases its internal energy, H <0). The increase in pressure causes shifting of the equilibrium towards less quantity of gaseous substances (as pressure is affected only by gaseous substances), and decrease in pressure shifts the equilibrium towards more quantity of gaseous substances. In case if quantities of gaseous substances among reactants and products are same, change in pressure does not cause shifting of chemical equilibrium. The increase in concentration of one of reactants causes shifting of equilibrium towards formation of products of reaction, and increase in concentration of one of reactional products shifts the equilibrium towards reactants. STRUCTURE OF ATOMS Atom is a complicated particle consisting of positively charged atomic nucleus and electronic shells where negatively charged electrons are located. The positive charge of a nucleus equals the sum of negative charges of electrons. The nucleus itself consists of positively charged protons and uncharged neutrons. Protons, neutrons and электроны carry the name «elementary particles». Using relative units, one can state that electrons are –1 charged, and protons are +1 charged. The mass of an atom expressed in nuclear mass units, is called a relative nuclear mass or mass number of an atom, Мr. It equals sum of masses of all elementary particles of the atom. As mass numbers of protons and neutrons are equal (1 a.m.u.), and masses of electrons are neglectably small (approximately 2,000 times less than corresponding masses of protons and neutrons) mass number of an atom equals sum of number of protons and neutrons. 21 Symbols of chemical elements are usually represented in the way: a X c b , where X is the symbol of the element; a is the mass number of the element; b is an element serial number in the Periodic Table (equals number of protons in the atom); c is the charge of the ion. Natural chemical elements exist in the form of a mixture of isotopes. Isotopes are atoms of the same chemical element with identical number of protons, but different mass numbers (number of neutrons). For 35 example, natural chlorine exists in the form of two isotopes: 17 37 17 Cl (17 Cl protons and 18 neutrons), and (17 protons and 20 neutrons). Relative masses of elements in periodic system of elements, are average masses of natural isotopes. Structure of electronic shells Electrons possess so called wawe dualism (simultaneously properties of a particle and a wave). In this connection, for the description of properties электрона enter special function which name state function of an electron or wave function, . It is entered in such a manner that the square of its module is proportional to probability to find out a particle (electron) in the given place at the appointed time (probability density). Wave function of an electron is called «orbital». It characterises energy and the form of spatial distribution of an electronic cloud. 22 Quantitative parities in the theory of a structure of atom are defined by the Shrodinger’s wave equation: h 2 2 2 2 U E 8 2 m x 2 y 2 z 2 where U is potential energy of the electron; Е is full energy of the electron; m is the mass of the electron; x, y, z are spacial co-ordinates of the electron; is the wave function; Consequence of the decision of Shrodinger’s equation Шредингера is the set of four quantum numbers which characterise behaviour of the electron in the atom. n, principle quantum number, it defines the general stock of energy of the electron, i.e.energetic shell. n = 1,2,3 … l, azimutal quantum number, defines the form of electronic orbital (subshell). l = 0,1,2 … (n-1). If l = 0 the orbital is called s-orbital (spheric movement of the electron). At l = 1 we have p-orbital (double-lobe movement form). Movement forms of d - and f - orbitals (l = 2 and 3 accordingly) have even more complicated character. The number of orbitals at an energetic level coincides with its number. So, for the first shell (n = 1) there is only one subshell (l = 0), that is 1s-orbital. Similarly for n = 2 (the second shelll) two subshells exist (l = 0, 1) or 2s, 2p-orbitals; for the third shell (n = 3, l = 0, 1, 2) 3s, 3p, and 3dorbitali exist etc. ml is the magnetic quantum number, it characterises projection of the magnetic moment of the electron on an external magnetic field, that is defines the spacial orientation of the electronic orbital. Its values are defined by azimutal quantum number: ml = l; (l-1); (l-2) … 0 23 For the orbital quantum number l = 0, magnetic quantum number has one possible value (ml = 0), that is only one way of orientation of sorbital in space is possible. Similarly we receive, that for p-orbitals (l = 1, ml =-1, 0, +1) there are three possible ways of orientation (along coordinate axes), for d-orbitals there are five possible ways of orientation (l = 2, ml =-2,-1, 0, +1, +2). To specify the concept electronic orbital, we can state that it represents a set of positions of electrons in the atom. Conditionally nuclear orbitals are designated in the form of cages (energetical cells): 1s 2s 2p 3s 3p 3d ms , the spin quantum number, defines the moment of spinning of the electron. As there are only two ways of spinning (clockwise and anticlockwise), the magnetic quantum number can accept two values: ms = l. Conditionally, electrons having different values of spin quantum number, are designated by opposite directed arrows: . Electronic formulae of atoms If the atom is in the ground state (does not possess superfluous energy) its electrons occupy the lowest energetic orbitals. Energy of an electron in multielectronic atoms depends not only on its attraction to a nucleus, but also from repultion from other electrons. Mutual influence leads to that energy of the electron depends not only on principle, but also on azimutal quantum number. 24 Klechkovsky’s rules 1. The increase in energy of electronic subshells goes as increase in the sum of the principle and azimutal quantum numbers (n+l). 2. In case of equality of the sum (n+l) the increase in energy of subshell goes as increase in the principle quantum number. Graphically it is possible to present the Klechkovsky’s rule in a kind: n 1 2 3 4 5 l 0 1s 2s 3s 4s 5s 1 2 3 2p 3p 4p 5p 3d 4d 5d 4f 5f Filling of orbitals by electrons occurs in a following order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p etc. Pauli's exclusive principle In one atom, there cannot be two electrons with an identical set of quantum numbers. Because everyone electronic orbital is characterised by a set of three quantum numbers (principle, azimutal and magnetic), electrons of the same orbital can differ only by the value of spin quantum number (ms = ). A consequence of a Pauli’s principle of is that one orbital can contain not more than two electrons. In connection with the previously mentioned at the first energetic shell, not more than two electrons can exist: Or 1s2; 1s And the second energetic shell can maximally contain 8 electrons: Or 2s22p6 etc. 2s 2p The maximum number of electrons at any shell N = 2n2, where n is the principle quantum number. 25 Hund’s rule In a subshell electrons fill orbitals so that the total spin quantum number becomes maximum (orbitals of a subshell are first filled by one electron each and only after all orbitals are filled, pairing of electrons takes place). For example, four electrons on a p-subshell can be arranged in two different ways: Or (ms) = + 1 (ms) = 0 As in the first case the total spin number is more,the first electronic structure is realized. Electronic formulae of atoms and ions The number of electrons in atom is defined by an element serial number in the Periodic table. Using the above rules and principles, for a sodium atom of (11 electrons) the following electronic formula is received: 2 2 6 1 11Na: 1s 2s 2p 3s 1s 2s 2p 3s The electronic formula of a Ti atom: 2 2 6 2 6 2 2 22Ti: 1s 2s 2p 3s 3p 4s 3d 1s 2s 4s 3d 2p 3s 3p If one electron do not suffice to full or half-full d-subshell (d10 or d -configurations), it is transmitted from the next s-subshell. As a result, the electronic formula of Cr atom looks like 24Cr: 1s22s22p63s23p64s13d5, 5 26 instead of 24Cr: 1s22s22p63s23p64s23d4, and atom of copper - 29Cu: 1s22s22p63s23p64s13d10, instead of 29Cu: 1s22s22p63s23p64s23d9. The number of electrons in a negatively charged ion - anion exceeds number of electrons in a neutral atom: 16S2 - 1s22s22p63s23p6 (18 electrons). At formation of a positively charged ion - cation – electrons are detouched first from the outer shell: 24Cr3 +: 1s22s22p63s23p64s03d3 (21 electrons). Electrons in atom can be divided into two types: internal and external (valence). Internal electrons occupy completely filled subshells, have low values of energy and do not participate in chemical transformations of elements. Valence electrons are all electrons of the outer electronic shell as far as electrons of not-filled subshells. Valence electrons take part in formation of chemical bonds. Unpaired electrons have special activity. The number of unpaired electrons defines the valence state of a chemical element. In case if not filled orbitals are available on the highest energetic level, unpairing of valence electrons may occur, and the valence state of the atom increases (formation of an excited state of the atom takes place). For example, valence electrons of sulfur are 3s23p4: 16S 3s 3p 3d In the ground state, the S atom has 2 unpaired electrons, so its valence state is II. At an expense of some energy one of paired electrons of sulfur can be translated on an empty d-orbital, that corresponds to the first excited state of the atom: 16S* 3s 3p 3d In this case, the S atom has four unpaired electrons, and its valence state equals IV. One of 3s- paired electrons can also be moved to a free 3d-orbital: 16S ** 27 3s 3p 3d In such a condition, the S atom of sulphur possesses the valence state VI. If the outer electronic shell has no free orbitals or subshells, unpairing of electrons is not possible and the atom can have onlt one valence state (example – O-atom): 8О 2s 2p THE PERIODIC LAW AND PERIODIC TABLE OF CHEMICAL ELEMENTS In 1869, the Russian chemist Dmitry Mendeleyev has shown that properties of simple substances, and forms and properties of chemical compounds of elements are in periodic dependence on nuclear scales of elements. As expression of this periodic law, the table, which reflects the law, was served. In 1914 the English scientist G.Mozli has shown, that the charge of a nucleus of an atom is numerically equal to an element serial number in the periodic table, so properties of elements and their compounds are in periodic dependence on a nucleus charge of the atom. The periodic table of elements reflects electronic structures of atoms. Each period (a horizontal series of periodic table) begins by an element in which electrons start occupying a new electronic shell with the principle quantum number that equals the number of the period). Groups (vertical columns) contain elements with identical number of valence electrons that equals the group number. Groups A contain selements (valence electrons occupy s-subshells). In case if valence electrons are on s - and p-subshells, they carry the name of p-elements. Elements with not fulfilled d - or f-subshells, are known as d - and felements. They occupy groups B of the periodic table. 28 Change of properties of chemical elements in the periods and groups of the Periodic table Chemical properties of elements are illustrated by interactions of their atoms. Properties of chemical elements can be divided into metallic (reducing, i.e. properties to lose electrons) and nonmetallic (oxidising, i.e. properties to gain electrons). Properties of chemical elements depend on strengths of attraction of valence electrons to a positively charged nucleui of atoms and are defined by following characteristics: Ionization energy (Ei) is an energy that is necessary for spending for a separation and removal an electron from atom, an ion or a molecule. Ionization energy is a measure of metallic (reducing) properties of elements: the lower the Ei, the stronger the metallic properties are. In groups at increase in a serial number of an element, the ijnization energy decreases, and in period - increases. Energy of electron affinity (Ea) is an energy that is allocated at joining an electron to an atom or a molecule. It characterises non-metallic (oxidising) properties of elements: the greater the value Ea, the stronger the non metallic properties are. In the periods from left to right energy of electron affinity and non-metallic (oxidising) properties of elements increase, and in groups from up to down they decrease. The half-sum of ionisation energy and energy of electron affinity is called electronegativity of atom. It increases with increase in non-metallic properties of elements. In the periodic table, non-metallic elements are settled down in groups A and occupy its right top part. Metallic elements of groups A are in the left bottom part of periodic table. All elements of groups B possess metallic properties. 29 CHEMICAL BONDS The structure of chemical compounds is defined by the nature of chemical bonds. The chemical bond arises at the interaction of atoms causing formation of chemically steady two-or multinuclear system (molecule, crystal, etc.). The formation of a chemical bond is connected with the general decrease in energy of a system of co-operating particles. The major characteristics of a chemical bond are bond energy, bond length, bond angles. Bond energy is a quantity of energy allocated at formation of a chemical bond. The more the bond energy, the stronger the molecule is. Bond length is a distance between nuclei of atoms in a molecule. Bond lengths are caused by sizes of reacting atoms and degree of overlap of their electronic shells. On a way of formation, three principal types of a chemical bond are distinguished. These are ionic, co-valent, and metallic. Covalent bond The chemical bond between the atoms carried out by shared electron pairs is called a covalent bond. It arises between identical atoms forming gaseous binuclear molecules, and also in the condensed state with participation of non-metallic atoms. There are two basic concepts of description of a co-valent bond: 1. Method of valence bond (VВ). 2. Method ща molecular orbitals (МО). 30 Both methods mutually supplement and do not exclude each other as they use various ways of approaches. Method of valence bonds According to the VB method, valency can be considered as number of formed shared electronic pairs. From the point of view of the exchange mechanism, valency of an element is defined by number of nonpaired electrons. Atoms can form a limited number of chemical bonds according to their valency. It corresponds to saturability of a covalent bond. Depending on number of unpaired electrons, atoms can form one, two or three co-valent bonds, i.e. a co-valent bond may be simple, double or triple. The strongest chemical bonds arise in a direction of a maximal overlap of atomic orbitals. As the orbitals have different spatial dispositions, therefore co-valent bonds are characterised by orientations. Depending on directions of overlapping, one can distinguish , and -bonds. -bonds are formed when two atomic orbitals overlap along an axis connecting nuclei of atoms. -bonds are formed when two atomic orbitals are site-overlapped. -bonds arise at overlapping of two d-orbitals, located in parallele planes. The hybridization process also influences orientation of a covalent bond. Hybridization is a mixing of of different subshells of an atom, electrons of which participate in formation of equivalent chemical bonds. Depending on hybridization type, hybrid orbitals have different position in space. 31 sp - linear, an interbond angle 180 sp2 - triangular, an interbond angle 120 sp3 – tetrahedral, an interbond angle 109 Shared electron pairs in a molecule are shifted to a more electronegative atom, thus a co-valent bond possesses a property of polarity. Molecules formed by identical atoms (Cl2, H2, etc.), have nonpolar bonds. The more the difference of electronegativities of the two atoms forming a chemical bond, the more polar it is. In case if an exchange mechanism of formation of a co-valent bond takes place, one of atoms (donor) delivers a pair of electrons, and another (acceptor) – a not-filled orbital. Therefore, a co-ordinate bond is formed. For example, formation of an ammonium ion (NH4 +) involves realization of both mechanisms of formation of a co-valent bond. The electronic formula of a nitrogen atom is 2s 2p It has five electrons on the valence shell 2 2 3 1s 7N 1s 2s 2p . of which three electrons are unpaired and form covalent bonds with three H-atoms (exchange mechanism). The lone electron pair of nitrogen participates in formation of a co-ordinate bond with an ion Н+ (nitrogen represents itself as the donor, and ion Н + as an acceptor of electrons). The method of valence bonds allows distinguishing concepts of valency and oxidation state. Valency of an atom characterises its ability to form co-valent chemical bonds. Oxidation state is a conditional charge on atom in a molecule if to assume, that shared electronic pairs are completely shifted to a more electronegative atom. 32 For example, valency of nitrogen in molecule NH3 is III (three covalent bonds). Since electronegativity of nitrogen exceeds that of hydrogen, all three formed shared electronic pairs are shifted towards nitrogen, giving to it a negative charge (oxidation state –3). The VB method ВС theoretically predicts structures and properties of many molecules and ions. Therefore, it cannot explain existence of some molecular ions (He2+, O22- ets). By means of this method, it is impossible to explain magnetic properties of some molecules, for example, О2 and В2. Method of molecular orbitals (МО) According to МО method, the molecule consists of a set of nuclei and electrons disposed on molecular orbitals. Main tenets of the MO method: 1. Each electron in a molecule occupies a certain energetic level (molecular orbital, MO) which is characterised by a molecular function and a corresponding set of quantum numbers. 2. The total number of formed МО equals the number of initial atomic orbitals. 3. Filling of MO occurs according th all the principles presented for atomic orbitals. 4. МО it is considered as a linear combination of atomic orbitals (MO – LCAO). Let us consider, for example, formation of molecule АВ. Valence electrons of each atom are on p-orbitals. If the wave function of the isolated atom A is А, and for the atom B it is В, so according to the МО method: АВ=С1АС2В 33 where С1 and С2 are coefficients considering the income of each atom in formation of a molecular orbital. . А E В (p*.) _ А В + А В (p) The new molecular orbital with a lower energy (p) is known as a binding orbital. As its energy is lower than the energy of an atomic orbital, electrons on it stabilize the molecule. The MO with a higher energy (p*) is called an antibinding orbital. Electrons on it tend to destroy the molecule. Stability of a molecule is described by a bond order, B.O. B.O. = ½ (No of binding electrons – No of antibinding electrons) If B.O. = 0 the number of electrons on binding orbitals equals the number of electrons on antibinding electrons. Such molecule is unstable and breaks up to initial atoms (does not exist). Conditions for formation of MO from AO are the following: close values of energies of overlapping atomic orbitals; considerable overlap of AO (formation of and -types of MO); identical spatial disposition of AO (рх - рх, instead of рх - рz). 34 On level power molecular орбитали two-nuclear molecules settles down in a following order: s(1s) s(1s) s(2s) s*(2s) x(2px) z(2pz) = y(2py) z*(2pz) = y*(2py) x*(2px) Examples of the description of molecules using the MO method (energetic diagrammes of molecules and molecular ions): 1. Molecule Н2 Energetic diagrammes shows transformation of atomic orbitals into molecular orbitals. e e e Е s* * 1s 1s s Н-atom Н2-molecule Each H-atom has only one s-orbital. While they interact, two molecular orbitals are formed. According to the principle of minimal energy, both two electrons in the H2 molecule occupy the sbinding orbital. 1 The s bond order is: B.O. = ½(2 – 0) = 1 This correlates with the single bond in H2 molecule from the point of view of the VB method. Н-atom 35 2. Molecule Не2 As each He-atom has two electrons paired on 1s-atomic orbitals, in the He2 molecule both binding and antibinding orbitals contain same number of electrons. B.O. = 0, so the molecule does not exist. Е s* * 1 s 1s s Не-atom 36 Не2-molecule Не-atom 3. Molecular ion Не2+ One can suppose that a molecular ion Не2+ can be formed if a Неatom interacts with the Не+-ion. Totally, 3 electrons are present in the spyce, two of which occupy a s-binding orbital and the rest one is disposited on the s*-antibinding orbital. B.O.= ½, i.e. the molecular ion Не2+ exists and forms a semi-bond from the point of view of the VBmethod. The existence of a Не+-ion was proved experimentally, aqnd it was found that the bond between two nuclei is twice weaker than that in the H2molecule. Е s* * 1 s 1s s Не-atom Не2+molecular ion Не+-ion 37 4 Molecule О2 Е p* p* 2p p p s* 2s 2s s O-atom 38 O2-molecule О-atom Electronic configuration of valence shell of Oatom is (2s22p4). The interaction of two sorbitals of two oxygen atoms is analogue to previous cases. p-Orbitals can form both and bonds. All 12 electrons of the two O-atoms occupy lower molecular orbitals. As two p* molecular orbitals are of same energy, O2 molecule has two unpaired electrons and possesses paramagnetic properties. Bond order B.O. = 2 which correlates with the double bond from the VB-method. Ionic bonds Ionic bond represents an electrostatic interaction between ions of opposite charges. Ionic bond can be considered as a limiting case of a polar co-valent bond where the difference of electronegativities of the two atoms forming a chemical bond exceeds 2). Usually it is considered that a ionic bond is formed at interaction of typical metals and typical non-metals. Energy of ionic bonds depends upon: 1. energy of electrostatic interaction between ions, i.e. it increases with increase in charges of ions and reduction of their radii; 2. energy of electronic affinity of non-metals which increases at increase in non-metallic properties of elements; 3. ionisation energy of atoms. Example: formation of a molecule of sodium chloride: Na + Cl NaCl Na Na + + e Еi = 495 kJ Cl + e Cl Еa = 345 kJ Na + + Cl - NaCl Ecolomb = 585 kJ Еbond = Еcolomb + Еa - Еi = 435 kJ Ionic bonds are not directed and not saturable. That defines a great stability of ionic crystals. INTERACTION OF MOLECULES (THE CONDENSED STATE OF SUBSTANCES) Chemical stability of molecules is shown only in systems, where distance between molecules is much more than their sizes (r10-9m). That corresponds to a gaseous state of a substance. 39 In case if the distance between molecules makes about 10-9 m (condencad state which may be liquid or solid) , arise forces of van der Waals which have electrostatic nature and are subdivided on: 1) orientational (dipole - dipole); 2) inductional (dipole - not polar molecule); 3) polarising (dispersive interaction of instantly induced dipoles of polarizable molecules). Hydrogen bonds have intermediate character between intermolecular interaction and a co-valent bond. It is a kind of interaction between positively polarised atom of hydrogen and negatively polarised atoms with high electronegativity (F, O, N, S, etc.). At the expense of the small size of an H-atom, it has ability to enter electronic shells of other atoms where there is an interaction that is intermediate between electrostatic interaction and covalent bond (interaction with lone electron pairs of non-metallic atoms). Hydrogen bond is indicated as Х - Н … Y (X, Y =F, O, N, S). CONCEPTS OF ADSORPTION Adsorption - is the takeover of one substance onto another surface. Absorption - is the takeover of one substance into another volume. Noncompensative forces of attraction and repulsion of molecules of a substance on the surface lead to the surface tension and the ability to adsorb molecules from the environment: 40 By the type of interaction between the molecules of the surface (adsorbent) and molecular environment (adsorbate), the adsorption can be classified as following: 1. Physical adsorption (only associated with intermolecular interactions. Such adsorption is reversible and is always accompanied by desorption); 2. Chemical adsorption which is accompanied by the chemical reactions on a surface, such as the occurrence of an oxide film on the metal surface. This adsorption is irreversible. The adsorption dependents on the temperature ( reduced by heating ) and the pressure ( increasing the pressure increases the adsorption of gas phase). The dependence of the adsorption of the adsorbate concentration (or pressure) at constant temperature is called adsorption isotherm: Г = f(P)T = Const или Г = f(С)T = Const According to the theory of Langmuir, adsorption occurs only in the places of maximum of attractive forces ( active sites ) . If all of the active surface sites are occupied, the further adsorption does not take place. If we assume that Г is the adsorption at any given time , and Г max is maximally possible adsorption (all active sites of the adsorbent are employed by adsorbate ) then Г / Гmax = (surface coverage); 1 – - the degree of free surface capable of adsorption. 41 The adsorption rate is proportional to the concentration of the adsorbate in the environment and the amount of free space on the surface of the adsorbent ( proportion of empty seats ), and the desorption rate is proportional to the number of seats occupied on the adsorbent surface : v(ads.) = k(ads.)( 1 – ) [adsorbate] v(des.) = k(des.) . Upon reaching the equilibrium, v(ads.) = v(des.); k(ads.)( 1 – ) [adsorbate] = k(des.) ; k(des.)/ k(ads.) = / ( 1 – ) [adsorbate] = К. Then, Г = Гmax K [adsorbate] / 1 + K [adsorbate] At the initial moment of adsorption, the adsorbate concentration on the adsorbent surface is small , ie 1 >> K [ adsorbate ] , the denominator is close to unity , Г = Гmax K [ adsorbate ] , ie, the adsorption is linearly dependent on the concentration of the adsorbate (line 1 in the graph) . After all of the active adsorption sites are occupied, the concentration of the adsorbate great 1 << K [ adsorbate ] , the denominator can be taken for K [ adsorbate ] , that is, Г = Г max (adsorption is constant and does not depend on the concentration of the adsorbate , which is reflected line 2 parallel to the axis of the abscissa ) . In between, the dependence of Г on the concentration of the adsorbate is nonlinear , which is reflected on the adsorption isotherm : Г 2 1 [adsorbate] 42 CONCEPT OF CATALYSIS The process of changing the rate of a chemical reaction due to the introduction of substances into the reaction system which is not part of the reaction products is called catalysis. Positive and negative catalysis are distinguished, ie the ones which accelerate or decelerate the rate of the reaction. Catalysts are substances that increase the rate of reaction and remain unchanged after it. Substances that slow down the rate of reaction are called inhibitors. Reactions in which one of the products is the catalyst of this process , called autocatalytic . The homogeneous and heterogeneous catalysis are distinguished. In the case of homogeneous catalysis, the reactants and the catalyst are in the same phase. An example is the oxidation of gaseous sulfur (IV) oxide to a sulfur (VI) oxide with nitric (IV) oxide as a catalyst. The phenomenon of homogeneous catalysis may be explained by the theory of intermediate compounds , according to which in the presence of a catalyst the reaction proceeds in several stages. Schematically, this may be expressed as : A + B = AB (reaction without catalyst is slow ) . In the presence of a catalyst : 1-st Stage A + K = AK ( fastly. AK - Intermediate ) 2nd Stage AA + B = AB + K ( Catalyst K after the reaction remains chemically unchanged ) . The main reason for accelerating action of homogeneous catalysts is to reduce the activation energy necessary for the reaction. In heterogeneous catalysis, the reactants and catalyst are in different phases (generally , the catalyst is a solid , the surface of which participates in the acceleration of the reaction) . An example is the reaction of oxidation of sulfur oxide (IV) with oxygen on the platinum surface . Heterogeneous catalysis is due to the adsorption of the reactants on the catalyst surface ( active areas ) . Increasing of the surface concentration increases the reaction rate. The reaction products are desorbed from the catalyst surface . 43 The rate of heterogeneous catalysis as the adsorption rate will be determined by the number of active sites on the catalyst surface , and a graph similar to the adsorption isotherm . The heterogeneous catalyst is always more active than the homogeneous catalyst (largely increases the reaction rate) . Biological ( enzymatic ) catalysis - a catalysis of biochemical reactions using biocatalysts - enzymes. Features of enzymatic catalysis: 1. The high catalytic activity of enzymes ( in hundreds times more active than the inorganic catalysts) . 2. Biocatalysts unlike inorganic catalysts have high specificity ( a single enzyme is generally catalyzes a biochemical reaction). 3. The need to create special conditions (even a small change in pH and temperature leads to a change in the catalytic properties of enzymes) . 44 SOLUTIONS Solutions are homogeneous systems of variable composition. Solutions consist at least of two components - solvent and solute. Solvents are accepted to be that substances which keep their aggregate states or which are of greater amounts. Amount (mass) of the solute in a mass or volume unit of the solution is nown as concentration of a solution. The most widespread concentration units of solutions are rhe following: Mass fraction represents a mass of substance in 100 g of a solution: m( solute) m( solution) (100%) Molar concentration (molarity) is a number of moles of a solute in one liter of a solution: CM ( solute) V( solution) m( solute) M ( solute) V( solution) Equivalent (normal) concentration is a number of equivalents of a solute in one liter of a solution: CE n( solute) V( solution) m( solute) M E ( solute) V( solution) Solubility is an ability of one substance to be dissolved in other under the set conditions. Quantitatively it is expressed by solubility factor, s. It equals concentration of the saturated solution under the given conditions. Solubility of substances depends on temperature and pressure: for liquid and solid solutes, it increases at rise in temperature, for gases - at fall of temperature and pressure increase. 45 Physical and chemical processes in solutions Interaction between molecules and ions of molecules of solute and solvent can consist of the several processes proceeding consistently or simultaneously. 1. Molecular dissociation of a solute: (АВ)k k AB 2. Interaction of molecules of solute with molecules of solvent (formation of solvates): AB + (n+m) S AB (n+m) S 3. Electrolytic dissociation (splitting of a solute into solvated ions): AB (n+m) S Ax + nS + Bx mS Substances which can form ions while being dissolved, are known as electrolytes. The quantitative characteristic of electrolytic dissociation is known as degree of dissociation: = Сdis / Сtot, where Сdis is the consentration of dissociated part of the electrolyte, and Сtot is its total concentration. According to the value of the degree of dissociation, the electrolytes can be devided into two groups: 1. Strong electrolytes (> 0.3 or 30 %). Among strong electrolytes there are some strong acids (HCl, H2SO4, HNO3, HClO4, HBr, HI), alkalis (soluble bases such as NaOH, KOH, Ca(OH)2, Ba(OH)2, etc.) and practically all salts. In solutions, strong electrolytes are practically completely broken up to ions (dissociation is irreversible and complete): 2 Al2 (SO4) 2 Al3 + + 3 SO4 46 2. Weak electrolytes ( 0.03 or 3 %). Among weak soluble electrolytes there are weak acids, ammonium hydroxyde NH4OH, and water itself. Dissociation of weak electrolytes is a reversible and stepwise process that is characterised by stepwise and overall equilibria constants (dissociation constants). For example, dissociation of phosphoric acid is a three-step process: [H ] [H2 PO4 ] H + + H2PO4 ; 1-st step: H3PO4 2-nd step: H2PO4 [H3 PO4 ] K1 = [H ] [H PO4 H + + HPO4 2 2 H + + PO4 3 [H2 PO4 ] K2 = 2 ; 2 ; [H ] [ PO4 3-rd step: HPO4 =8103 K3 = 3 [ HPO4 ] ] =6108 ] =21012 Overall process: [ H ]3 [ PO4 H3PO4 3H + + PO4 3 ; K= 2 [H3 PO4 ] 3 3 ] = К1К2К3=11021 where [H+], [H2PO4 ], [HPO4 ], [PO4 ], [H3PO4] are equilibrium concentrations of ions; К1, К2, К3 – stepwise dissociation constants; and K is the overall dissociarion constant. 47 Reactions of ionic exchange Reactions of ionic exchange in solutions occur between ions of strong electrolytes and molecules of weak electrolytes and insoluble substances. They proceed towards formation of precipitates, gases, and molecules of weak electrolytes. Na2SO4 + 2HNO2 2NaNO2 + H2SO4 soluble strong soluble weak (reaction in a molecular form) 2Na + + 2NO2 2Na + + SO42 + 2H + + SO42 (full ionic form of the reaction) 2NO2 + 2HNO2 2HNO2 + 2H + (net ionic form) The character of the dissociation defines the properties of chemical compounds in solutions: HCl H + + Cl (acids form hydrogen ions Н + while dissociation); NaOH Na + + OH NaCl Na + + Cl (bases dissociate to produce ions of hydroxide OH ); (salts form metallic cations and anions of acids). The main reaction that reflects acidic and basic properties is the reaction of neutralization (acids interact with bases to produce salts): Na2SO4 + 2H2O 2NaOH + H2SO4 2Na + + 2OH 48 2Na + + SO42 + 2H2O 2H2O 2OH + 2H + + 2H + + SO42 There are electrolits that can participate in chemical reactions both as bases and as acids. Such electrolytes are called amphoteric. Among them there are Zn(OH)2, Pb(OH)2, Sn(OH)2, Be(OH)2, Al(OH)3, Cr(OH)3, As(OH)3, and some others. These substances are capable to react both with acids and with the bases, forming salts as products of reaction of: Al(OH)3 + 3HCl AlCl3 + 3H2O Al(OH)3 + 3H+ Al3+ + 3H2O Al(OH)3 + 3NaOH Na3 [Al(OH)6] Al(OH)3 + 3OH- [Al(OH)6]3Sn(OH)2 + 2HCl SnCl2 + 2H2O Sn(OH)2 + 2H+ Sn2+ + 2H2O Sn(OH)2 + 2NaOH Na2 [Sn(OH)4] Sn(OH)2 + 2OH- [Sn(OH)4]2- DISSOCIATION OF STRONG ELECTROLYTES In solutions of strong electrolits owing to their full dissociation, concentration of ions is great; therefore, properties of such solutions will depend on degree of interaction of ions as with each other, and with polar molecules of solvent. So, concentreations of ions are replaced by their activities. Activity is a visual concentration of an ion involving its interaction with other ions of the solution: a = fC (f - activity factor). If f = 1 ions are free and do not co-operate among themselves (a=C). If f <1 ions co-operate (a <C). The less the activity factor, the more interaction between ions exists in the solution. 49 The activity factor depends on total concentration of all the ions in a solution (ionic strength of a solution): = l Ci Zi2, where - ionic strength; Ci - concentration of ions in a solution; Zi charges of ions. log f 0.5Z 2 1 (Debaue-Huckel’s equation) For diluted solutions of strong electrolytes with <<1, log f 0.5Z 2 DISSOCIATION OF WEAK ELECTROLYTES While dissociation of weak electrolytes takes place, equilibrium is established. CH3COOH CH3COO + H + Thus, if the total concentration of the electrolyte equals C, and degree of its dissociation is , Cdis = ∙ C [CH3COO-] = [H +] = C [CH3COO-] = C – C K H CH COO 3 CH 3COOH For 1 K=C2 50 C 2 2 C 2 2 C 2 C C C 1 1 and K C The resulted equation expresses the Ostwald’s dilution law according to which degree of dissociation of a weak electrolyte increases with dilution of a solution. Addition of common ions in a solution of a weak electrolyte causes shifting of equilibrium of the reaction towards reduction of dissociation (effect of a common ion). ELECTROLYTIC DISSOCIATION OF WATER Water is a weak electrolyte that dissociates according to the equation: Н2О Н + + OH. o At the temperature of 22 C, the equilibrium of the dissociation process establishes such that: [H +] [OH ] = KH2O = 1014 (ionic product of water). In a neutral solution [H +] = [OH-] = In an acidic solution [Н +]> [OH-]; In an alkaline solution [H +] <[OH-]; 1014 = 107 mol/l. [H +]> 107 mol/l. [H +] <107 mol/l. Knowing concentration of one of the ions, for example [Н +] and ionic product of water, it is possible to calculate concentration of another type of ions [OH-]. The negative logarithm of concentration of of hydrogen ions (or the negative logarithm of activity of ions of hydrogen) is named рН: рН = – log [H +] In neutral solutions at 22оС рН = 7. In acidic solutions рН < 7. In alkaline solutions рН > 7. 51 Acid-base indicators are substances, changing colouring in certain area of value pH a solution. Weak organic acids or bases, which molecules and ions have different colouring can be indicators. Area of transition of colouring of some indicators The indicator Methylorange Phenolphtalein Litmus Colour Area of transition of colouring, рН acidic form red colourless. red alkaline form yellow red dark blue 3.2 – 4.5 8.2 – 10.0 6.0– 9.,0 EQUILIBRIA IN SOLUTIONS WITH PRECIPITATES In the saturated solution of a sparingly soluble electrolyte, dynamic equilibrium between a firm phase and a solution is established which can be presented by the equation: CaCO3 (s) Ca2 + + CO32 For this equation, using the law of mass action, we can write down the expression of equilibrium constant: [Ca 2 ][CO3 2 ] [CaCO3 ] Ks = where [Ca2 +], [CO32 ] are equilibrium concentrations of ions in the solution; [CaCO3] is the concentration of the electrolyte in a precipitate (in a firm phase), it is constant. 52 Having increased Ks on a constant [CaCO3], we will receive a constant named solubility product, Ksp:. Ks [CaCO3] = Ksp = [Ca2 +] [CO32 ] Solubility product is a product of concentration of ions of a sparingly soluble electrolyte in its saturated solution in their degrees of stoichiometric factors. Numerical values of the solubility product of sparingly solyble electrolytes are presented in special help tables. In the presence of the common ions, the eqiolibrium is displaced towards deposit formation (effect of the common ion). In the presence of the strong electrolytes with no common ions, the mobility of ions in the solution decreases, and the equilibrium is shifted towards an increase in the precipitate dissolution (salt effect). At mixing of the solutions containing ions giving an insoluble salt, the precipitation starts. During the first moments of time, the concentrations of ions in a solution are high, and the precipitation takes place. Within time, gradually concentrations of ions decrease and an equilibrium is established. Condition for formation of a precipitate: the product of concentration of ions in the solution exceeds the value of solubility product of the given compound. For example, [Ca2 +] [CO32 ]> Ksp (CaCO3). Condition for dissolution of a precipitate: the product of concentrations of ions in the solution deceeds the value of the solubility product of the given compound ([Ca2 +] [CO32 ] <Ksp (CaCO3)). 53 Dependence of the solubility product and solubility of sparingly soluble electrolytes Consider the dissolution scheme of a sparingly soluble electrolyte of the KxAy type as KxAy x Ky + + y Ax The expression of the solubility product for this equation looks like: Ksp = [Ky +] x [Ax ] y Let us designate the molar solubility of the electrolyte as “S”. Then the solution will contain cations in the concentration [Ky +] = xS (mol/l), ans anions in the concentration [Ax ] = yS (mol/l). If to substitute these designations in expression of the solubility product: Ksp = [xS] x [yS] y, From here, we can find the solubility S: s x y K sp xxy y BUFFER SOLUTIONS Buffers are used to maintain a constant pH value in the sample solution upon addition thereto of small amounts of strong acids, strong bases, or dilution. As the buffer solutions, a mixture solution of weak acids or weak bases and their salts or mixtures of salts of polybasic acids with different degrees of substitution usually used. The following table shows examples of the most commonly used buffers and pH, they support: 54 Composition of the buffer solution A mixture of CH3COOH andСН3СООNa The mixture NаН2РО4 and Nа2НРО4 A mixture of NH4OH and NН4С1 The name of the buffer acetate buffer рН phosphate buffer 6.5 ammonia buffer 9.25 4.7 Buffer systems can bind both H + and ОН ions at the addition of strong acids and bases to produce weak electrolytes that slightly alterin the pH of the solution. Example: Acetate buffer solution comprises a mixture of CH3COOH and CH3COONa. The dissociation of a weak electrolyte - acetic acid – is reflected by the reaction equation: CH3COOH CH3COO + H+ and is described by the equilibrium constant: K a= [CH 3 COO ] [ H ] = 1.8 10 5 [CH3 COOH ] By adding sodium acetate, the concentration of CH3COO ions increases and is determined by the concentration of the salt: [CH3COO ] Csalt. The dissociation of a weak electrolyte is reduced by the introduction of a common ion, so [CH3COOH] Cacid. Thus, C K a= [H ] salt ; C acid 55 [H+] = Ka C acid ; C salt C pH = –log[H+] = pKa – log ( acid ), C salt where pKa = - log Ka. Thus, pH of a a buffer solution is not dependent on the concentrations of the components but is determined by their mole ratio. By adding small amounts of a strong acid or a strong base (alkali), the components of the buffer solution translating them into weak electrolytes: Examples. 1. At addition of NaOH, it can react with the qacetic acid: CH3COOH + NaOH = CH3COONa + H2O (the salt concentration increases accordantly to the concentration of the added amount of the alkali, and the concentration of the weak acid decreases to the same value): pH = –log[H+] = pKa - log ( С acid NaOH ). C C salt NaOH C 2. The addition of a strong acid provokes its interaction with the sodium acetate: CH3COONa + HCl = CH3COOH + NaCl, C pH = –log[H+] = pKa - log ( acid С NaOH ). C C salt NaOH Since the change of the ratio of the concentrations is much less than their sum or the difference, the total the pH value changes insignificantly. 56 The amount of a strong acid or a strong base to be added to the buffer solution to change the pH of one liter of a solution thereof to a unit is called a buffering capacity (B). It can be calculated with respect to the acid (Ba) or base (Bb). C a Va ( pH1 pH 2 ) V solution C V b b Bb= ( pH 2 pH1 ) V solution Ba= wherein Ba and Bb - buffering capacity for the acid and base respectivelyy; Ca and Cb - concentration of added acid and base; pH1 and pH2 - initial and final pH of the solution; Va and Vb – volumes of the added strong acids and bases. DIRECTION OF REACTIONS OF IONIC EXCHANGE Reactions of ionic exchange are irreversible only in that case when weak electrolyites or sparingly solyble electrolytes exist among the products. In that case, when weak electrolytes and precipitates are present both among the reactants, and among the products, the reactions of ionic exchange are reversible and are shifted either towards the products (the reaction is possible to proceed) or towards the reactants (the proceeding of the reaction is not possible). For determination of the possibility of proceedings of such reactions, the equilibrium constant must be calculated. Rule for calculation of the equilibrium constant. To calculate an equilibrium constant of the reaction, write down the product of all the 57 possible constants for the compounds among the products of the reaction in the numerator of the fraction expressing equilibrium constant (by the possible constants, the solubility products of sparingly soluble electrolytes, dissociation constants of weak electrolytes - acids or the bases, instability constants of of complex ions, ionic product of water are supposed), as well as the same for reaction products should be added to the fraction denominator, ranking all the constants in the degrees corresponding to the stoichiometric factors of the corresponding substance in the equation of the reaction. The multibasic acids are represented by their overall dissociation constants which are products of their stepwise constants. In the case if the equilibrium constant of the reaction, К(reaction)> 1, the equilibrium of the reaction is shifted towards the products, and the reaction of ionic exchange is possible. In the case if К(reaction)<105, the equilibrium of the reaction is practically completely shifted towards the reactants, and the reaction proceeding is impossible. In tha case if the К(reaction) is concluded in limits from 1 to 105, the reaction proceeding becomes possible in excess of one of the reagents. Example. Let us determine the possibility of dissolution of CaCO3 in the acetic acid. The molecular reaction of the process can be written as: СaCO3 + 2 CH3COOH Ca (CH3COO) 2 + H2CO3 In the net ionic form the reaction is presented as following: СaCO3 + 2 CH3COOH 58 Ca2 + + 2 CH3COO - + H2CO3 The equilibrium constant of the reaction may be calculated according to the formula and using the corresponding table data: Ksp(CaCO3 ) K 2 (CH 3COOH ) K (reaction) 6.6 10 2 . 1 2 K ( H 2 CO3 ) K ( H 2 CO3 ) The value of the К(reaction) specifies that calcium carbonate can be dissolved in excess of acetic acid. HYDROLYSIS OF SALTS If the salt dissolved in water contains ions-rests of weak acids or the weak bases there is a process of hydrolysis of salt - exchange reaction of ions of salt to molecules of the water leading to formation of molecules and ions of new weak electrolits. Key rules for writing the reactions of hydrolysis: 1. Only anions of weak acids and cations of weak bases which are a part of the salt are exposed to hydrolysis process. 2. Hydrolysis is a stepwise process. At each step, one hydrolyzed ion reacts with one molecule of water. 3. Under common conditions, hydrolysis proceeds only on the first step. Hydrolysis amplifies at heating and dilution of the solutions of the salts. 4. As a rule, hydrolysis is a reversible process. Its equilibrium it is possible to be displaced by addition of small amounts of strong acids and alkalis. 59 Types of hydrolysis reactions. 1. Salt is formed by the ions of a strong base and a strong acid (for example, NaCl, KNO3, etc.). NaCl + H2O hydrolysis does not proceed (NaOH is a strong base, HCl is a strong acid). 2. Salt is formed by the ions of a strong base and a weak acid (for example, Na2CO3, KSCN, etc.). Na2CO3 + Н2О anions are involved in hydrolysis (NaOH is a strong base, H2CO3 is a weak acid). CO32 + HOH Na2CO3 + HOH HCO3 + OH (alkaline medium, рН> 7). NaHCO3 + NaOH (1 step of hydrolysis). Addition of alkalis (NaOH) to the above solution (OH are formed within hydrolysis), causes depression in hydrolysis. Countrary, addition of acids strengthens hydrolysis at the expense of reaction Н ++ OH Н2О in which result concentration of ions OH in a solution decreases, and balance of hydrolysis is shifted towards the products. Hydrolysis amplifies and the second step becomes possible: HCO3 + HOH NaНCO3 + HOH H2CO3 + OH H2CO3 + NaOH (2-nd step of hydrolysis). 3. Salt is formed by the ions of a weak base and a strong acid (for example, AlCl3, FeSO4, etc.). AlCl3 + H2O cations are involved in hydrolysis (Al(OH)3 is a weak base, НCl is a strong acid). Al3 + + HOH AlOH2 + + H + (alkaline medium, рН <7) AlCl3 + HOH AlOHCl2 + HCl (1-st step of hydrolysis). 4. Salt is formed by the ions of a weak base and aweak acid: The salt is water soluble (for example, (NH4)2CO3, NH4NO2, etc.). 60 (NH4)2CO3 + H2O both cations and anions are involved in the hydrolysis process: 2NH4 + + CO32 + HOH NH4OH + HCO3 + NH4 + (рН 7) (NH4) 2CO3 + HOH NH4OH + NH4HCO3 The salt is insoluble in water (for example, FeS, ZnSiO3, etc.). FeS + H2O insoluble salts do not undergo hydrolysis. Some salts are irreversibly destroyed by water (for example, Al2S3; Al2(CO3)3; Cr2S3; Cr2(CO3)3; Fe2S3; Fe2(CO3)3): Fe2S3+6H2O 2Fe (OH) 3 +3H2S Calculations of pH values of the solutions of the salts Quantitatively hydrolysis is described by hydrolysis constant (Kh) and degree of hydrolysis (h). NaCN + H2O NaOH + HCN CN - + H2O OH - + HCN; pH> 7 It is possible to express a hydrolysis constant through concentration of ions in a solution taking into account that concentration of water practically does not change, or by a rule of calculation of an equilibrium constant of a reversible reaction of hydrolysis: Kh= KH O [OH ][ HCN ] 2 K [CN ] acid NH4Cl + H2O NH4OH + HCl NH4 + + H2O H + + NH4OH; pH <7 61 [ H ][ NH 4OH ] K H 2O Kh = K base [ NH 4 ] NH4CH3COO + H2O NH4 + + CH3COO Kh = NH4OH + CH3COOH + H2O NH4OH + CH3COOH KH O [ NH 4OH ][CH 3COOH ] 2 K [ NH 4 ][CH 3COO] acid K base Degree of hydrolysis “h” defines that part of salt, which has undergone to hydrolysis under the given conditions: h С hydrolyzed . С total. Hydrolysis degree is related to a hydrolysis constant: Kh = h 2C . 1 h In the case when h <<1, Kh = h2C. Example. Calculation of the рН of a 0.1 M solution of potassium phosphate. Let us consider, that the hydrolysis basically proceeds on the first step: PO43 Kh = 62 + HOH HPO42 + OH KH O 10 14 2 = = 7.69x10-3 12 III 1,310 K H PO 3 4 h= Kh C K PO 3 4 = 7.6910 3 = 0 .1 7.69 10 2 = 2.77 10-1 [OH ] = hC = 2.77 10-1 * 0.1 = 2.77x 10-2; [H +] = 10-14 / [OH ] = 10-14 / 2.77 10-2 = 3.61 10-13; рН = -log [H +] = 12.44. Irreversible hydrolysis The equilibrium of hydrolysis reactions in some cases may be irreversibly shifted towards the products iif a product of one of hydrolysis steps will be sparingly soluble: Bi (NO3) 3 + HOH BiOH (NO3) 2 + HNO3 Bi3 + + HOH BiOH2 + + H + At dilution of the solution, the equilibrium of the reaction is shifted to the right at the expense of formation of a precipitate of bithmuth oxonitrate: BiOH(NO3) 2 + HOH Bi(OH) 2NO3 + HNO3 BiONO3 + H2O BiOH2 + + HOH + NO3 BiONO3 + H + The hydrolysis of antimony(III) chloride proceeds similarly. 63 Mutual hydrolysis Ions Н + (or OH ) can combine together into water molecules in the case when two salts with different types of hydrolysis are mixed together. That will cause mutual strengthening of hydrolysis of both salts and as a result - formation of products of hydrolysis (mutual hydrolysis). For example, at mixing of solutions Na2CO3 and AlCl3 in which accordingly there is a surplus of i OH and Н + ons accordingly, mutual strengthening of hydrolysis leads to allocation СО2 and formation of a precipitate of Al(OH)3: 2AlCl3 + 3Na2CO3 + 3H2O 2Al(OH)3 + 3CO2 + 6NaCl 2Al3 + + 3CO32 + 3H2O 2Al(OH) 3 + 3CO2 In other similar cases, a least soluble of possible products of hydrolysis is formed while a mutual hydrolysis. For example, the solubility of a copper(II) hydroxo carbonate (CuOH)2CO3 is less than the one of copper hydroxide Cu (OH)2. Therefore, mixing of solutions of CuSO4 and Na2CO3 will lead to formation of (CuOH)2CO3: 2CuSO4 + 2Na2CO3 + H2O (CuOH)2CO3 + CO2 + 2Na2SO4 2Cu2 + + 2CO32 + H2O (CuOH)2CO3 + CO2 Mutual hydrolysis is an irreversible process. COLLOIDAL SOLUTIONS Colloidal solutions are ultra-microheterogenous systems consisting of two phases: a dispersion medium and a dispersed phase. The dispersed phase comprises colloidal solutions of molecular aggregates of a size of 10-7 – 10-5 cm. 64 Classification of colloidal solutions 1. Physical state of the dispersed phase and the dispersion medium. Dispersed phase gas liquid solid Dispersion medium gas liquid solid aerosols aerosols (clouds, fog) (smoke, dust) emulsions Sols, (milk, oil) suspensions (natural waters) solid emulsion gels Litosols (capillary system, (soil, pearls, (minerals, activated carbon) cheese) alloys) emulsions (foam) 2. Interaction of the dispersed phase and the dispersion medium (the type classification is used only for the liquid dispersion media) . Colloidal systems are divided into lyophobic and lyophilic . In the lyophobic colloidal systems, the dispersed phase weakly interacts with the dispersion medium . Such systems can not be obtained spontaneously dispersed phase always represent molecular aggregates (irreversible colloids) . In the lyophilic colloid systems (reversible colloids), the dispersed phase is soluble in the dispersion medium . Its particles are individual molecules of large size (macromolecules). 65 Properties of lyophobic colloids 1. Heterogeneous system (the interface of the dispersed phase and dispersion medium). 2. Colligative properties of the colloid practically do not differ from those of the dispersion medium. 3. Cannot be divided by mechanical methods (sedimentation , filtration, etc.) . 4. Poorly pass through semipermeable membranes and films. 5. When light passes through colloid solutions, a Tyndall’s cone is observed deu to diffraction of light on the colloidal particles. 6. Since the shortwave light diffraction occurs to a greater degree than the diffraction wavelength radiation , the colorless colloidal solutions are blue in reflected light, and red in the transmitted light. 7. The colloidal particles are charged and can move under the influence of an electric field (electrophoresis). Colloidal micelle structure and electrical properties of colloidal systems Colloidal particles are small in size and can be suspended indefinitely. This determines the kinetic stability of colloidal systems . On the other hand, a large colloidal particle surface defines excessive surface tension and the tendency to "sticking" of colloidal particles, which reduces the energetic stability of colloidal systems. Stabilization of colloidal systems is due to the adsorption of the dispersion medium or the electrolyte ions present in the solution. On the border of the colloidal particles, the electric double layer is formed. Thus. colloidal particles acquire the same charge , which prevents "sticking". Ions that form the charge of colloidal particles are called potential forming ions. The- counterions of the opposite charge are located in the adsorption and diffusion layers of the micelle. 66 The ion exchange occurs between the adsorption and diffusion layers. The colloidal particles are charged and the micelle electrically neutral. Example. The structure of the micelle of the colloidal solution of silver iodide . The formation of silver iodide can be represented as : KI (excess) + AgNO3 AgI + KNO3 The micelle structure : {m[AgI] nI (n-x)K+} x xK+ where "m" is the number of molecules of silver iodide in the core; "N" is the number of anions in the first adsorption layer of the colloidal particle ; n + x = m; typically m >> n. If there is an excess of silver nitrate in the solution, the silver ions are primarily grouped on the surface of the colloidal core constructing additional crystal lattice of the solid phase and nitrate anions serve as counterions. On the surface of the micelle core an electric potential of the interphase is generated ( - potential), which decreases with increasing distance from the core surface by adsorption of counterions. If placed colloidal particle in an electric field, the counterdiffusion layer is detached from the colloidal particles, and moved in the direction of the corresponding electrode, and a colloidal particle - in the opposite direction. The absorption and diffusion layers are devided by a sliding surface). The potential on the sliding surface is called electrokinetic potential or ( zeta ) potentials . - potential is a measure of the stability of colloidal systems : if < 30 mV, coagulation (destruction) of the colloidal particles occurs. Introduction of indifferent strong electrolytes into a colloidal solution decreases the -potential which is related to the increase in the concentration of counterions and contraction of the diffuse layer. In the case where the concentration of the electrolyte is sufficiently high, the isoelectric state ( = 0) can be achieved. 67 At introduction of a strong common electrolyte, there is an additional adsorption of the potential forming ions on the surface of colloidal particles , which leads to an increase of the - potential. After reaching the maximum possible adsorption of a common electrolyte, the latter starts to act as the indifferent electrolyte, ie the - capacity decreases and the coagulation of colloidal particles occurs. Coagulating ability of an electrolyte is characterized by the coagulation threshold - the minimum concentration of the electrolyte , which causes coagulation of the colloidal solution. Coagulation threshold decreases with increasing the charge of the coagulating ion (the ion charge sign is the same that the charge of the counterions of the colloidal particle). For the mono- , di-, and triply charged ions, the ratio of their coagulation threshold is 1:11:72 . In some cases, the coagulation process is reversible. The process of reverse transformation from coagulate to colloidal solution is called peptization or disaggregation. The stability of lyophobic sols increased by adding small amounts of solutions of high molecular compounds (e.g., gelatin solution, tannin et al.). The protective action of solutions of high molecular compounds is due to the formation of the protective layer which is adsorbed on the surface of the colloidal particle. The characteristics of the protective action is protective number - minimum amount ( mg ) of the solid substance that prevents coagulation of 10 ml of the colloidal solution by adding a strong electrolyte in an amount determined by the threshold of coagulation. Properties lyophilic colloids (solutions of highmolecular compounds , HMC) 1. Homogeneous systems. 2. They are formed spontaneously and are thermodynamically stable . 3. Colligative properties are determined by the concentration of the HMC. 68 4. The colloidal particles can be charged (ionic HMC) or be uncharged. 5. Coagulation threshold of the electrolytes significantly exceeds the one for lyophobic colloid systems . The mechanism of action is associated with coalescers desolvation (salting) of HMC. In some cases, the drops form a second liquid phase - structured liquid approaching the jelly by properties. This phenomenon is called coacervation and is typical for many proteins. The concentration of the HMC in coacervate drops increases , but otherwise the solution - is reduced compared with the original. 6. Concentrated solutions of HMCare characterized by high viscosity. Ways of preparation of colloidal solutions 1. The dispersion methods (grinding of large aggregates of molecules up to the size of the colloidal particles): a) mechanical dispersion (using mechanical devices - ball mills , grinders , etc.); b) electrical dispersion (grinding with an electric current) is used for the preparation of sols of metals and other conductive materials; b) peptization method (the precipitate is washed with freshly prepared stabilizer electrolyte solution . EXAMPLE . Peptization of potassium hexacyanoferrate(II) by the oxalic acid solution: K4[Fe(CN)6] + FeCl3 KFe[Fe(CN)6] + 3KCl H2C2O4 2H+ + C2O42 {m[KFe[Fe(CN)6]] n C2O42 (n-x)H+} x xH+. 2. The condensation methods (formation of aggregates of individual molecules, i.e. the formation of colloidal solutions in the course of chemical reactions): a) an ionic exchange reaction 3H2S (excess) + As2O3 As2S3 + 3H2O 69 {m[As2S3] nS2 (n-2x)H+} x 2xH+ b) redox reactions: 2HAuCl4 + 5K2CO3 2KAuO2 + 5CO2 + 8HCl + H2O KAuO2 + HCHO Au + HCOOK + H2O {m[Au] nAuO2 (n-x)K+} x xK+ 2H2S + O2 S + 2H2O 2H2S + 7O2 2H2O + H2S5O6 {m[S] nS5O62 (n-2x)H+} x 2xH+ c) hydrolysis reactions: FeCl3 + H2O FeOHCl2 + HCl FeOHCl2 + H2O Fe(OH)2Cl + HCl Fe(OH)2Cl FeOCl + H2O Fe(OH)2Cl + H2O Fe(OH)3 + HCl {m[Fe(OH)3] nFeO+ (n-x)Cl}X+ xCl {m[Fe(OH)3] nH+ (n-x)Cl}X+ xCl OXIDATION-REDUCTION REACTIONS (REDOX REACTIONS) Oxidation-reduction are those which proceed with the transmission of electrons from one particles (atoms, molecules or ions) to another therefore the oxidation states of some elements in the given particles change. Oxidation state is a conditional charge of an atom in a molecule provided that all polar bonds are considered as ionic. 70 Oxidation process is a process of gaining of electrons; it is accompanied by increase in the oxidation state of an element. Reduction process – is a process of gaining of electrons; it is accompanied by reduction of the oxidation state of an element. The oxidizer is an ion or a molecule which is taking away electons from other elements of an ion or molecule. In the course of oxidationreduction reaction, the oxidizer reduces. Reducer is an element, an ion or a molecule, providing electrons to other elements, ions or molecules. In the course of oxidation-reduction reaction, the reducer oxidises. Typical oxidizers: active nonmetals (O2; F2; Cl2), concentrated solutions of H2SO4; solutions of HNO3 of any concentration; KMnO4; K2Cr2O7 and other substances containing elements in their highest oxidation states. Typical reducers: active metals; H2; C; CO; H2S; HI and other substances containing elements in the minimal (negative) oxidation states. Elements in intermediate oxidation state possess redox duality, i.e. can show both properties of oxidizers, and properties of reducers (Na2SO3, KNO2, etc.). In oxidation-reduction reactions the following ions usually do not change their oxidation staes: ions of alkaline and akline earh metals; hydrogen ions (except reactions of metals with acids-non oxdizers); oxygen ions (except reactions with participation of hydrogen peroxide, H2O2); the ions forming molecules of acids or alkalis, defining reaction of the environment. In oxidation-reduction reactions, the following changes of oxidation states of elements are usually observed: Mn+2 7 K MnO4 4 Mn O2 (In the acidic environment, for example, MnSO4, MnCl2, etc.) (In the neutral environment) 71 6 K 2 MnO4 4 (In reactions with the metals of low activity, and non-metals) S0 (In reactions with the metals of average activity) S O2 6 H 2 S O4 (In the alkaline environment) concenhtrated 2 H2 S (In reactions with active metals) 4 (Heavy metals and non-metals) N O2 5 H N O3 concentrated . 1 N2 O . 2 NO (Active metals) (Heavy metals and non-metals) 5 H N O3 diluted . 0 N2 . (Active metals) (Very diluted nitric acid with very active metals gives NH4NO3 as a product of reduction (nitrogen in the oxidation state-3). 72 6 K 2 Cr2 O7 Cr 3 OH H OH H 6 3 3 4 Na2 S O3 Mn 2 Mn 2 5 Na Bi O3 1 H2 O 2 1 H2 O 2 Cl Sn+2 Pb+2 In the alkaline environment (for example, Na3 [Cr (OH)6]) 5 (In reactions with K N O3 oxidizers) 6 (In reactions with Na2 S O4 oxidizers) 7 (In the acidic H MnO4 environment) 4 (In neutral and alkaline Mn O2 environments) (In the acidic Bi 3 environment, for example, Bi (NO3) 3) 2 (In reactions with H2 O reducers) 0 (In reactions with O2 oxidizers) 0 (In reactions with strong Cl 2 oxidizers) Sn+4 Pb+4 [Cr ( OH )6 ]3 K 2 Cr O4 K N O2 In the acidic environment (e.g., CrCl3) Direction of oxidation-reduction reaction Oxidation-reduction properties of substance define on their oxidising ability, числено expressed through redoks - potential . The real potential pays off on the Nernst’s equation: 0,059 = 0+ n log (Cox / C red), 73 Where 0 - standard electrode potential of the given process, n - quantity of accepted electrons, Cox and C red - concentrations of the oxidised and reduced forms accordingly. The more value , the is stronger oxidising properties of the oxidised form of the given compound. The direction of the oxidation-reduction reaction is defined on the electromotive force (E): E = (ox) -(red) If E> 0, the direct reaction is possible. If E <0, the direct reaction is impossible (reaction goes to the opposite direction). Standard electromotive force E0 is connected with standard Gibbs’energy G 0 nFE On the other hand, G the reaction K: 0 0 = - G 0 defines on the equilibrium constant of 0 G = - 2.3RT log K Thus, nFE 0 = 2.3RT log K 0 At the temperature of 25 C (298 K), the equation becomes: nE 0 log K = 0.059 . Classification of redox reactions All oxidation-reduction reactions can be divided into three groups: 1). Reactions of intermolecular oxidation-reduction are reactions in which the exchange of the electrons occurs between the atoms of different molecules: Fe + CuSO4 = FeSO4 + Cu. 74 2). Reactions of intramolecular oxidation-reduction- in such reactions an oxidizer and a reducer are in the same substance: 2КClO3 = 2KCl + 3O2. 3). Reactions of disproportionation are those in which the same atoms in molecules co-operate with each other as an oxidizer and a reducer because these atoms have intermediate oxidation states: 3К2MnO4 + 2H2O = 2KMnO4 + MnO2 + 4KOH. Composition of the equations of redox reactions The method of half-reactions is based on drawing up of ionic equations for processes of oxidation and reduction with the account of рН environments of the given reaction. Strong electrolytes in the given method are registered in the form of ions, and the weak electrolytes- in the form of molecules. The ionic scheme includes ions and molecules showing oxidation-reduction properties, and the ions characterising the environment (Н + and water molecules in the acidic solutions, OH ions and water molecules in the alkaline environment, and all of the above in the neutral environment): Ca + HNO3 (разб) NH4NO3 +... Ca0 - 2 e Ca2 + ( 4) NO3 + 10H + + 8 e NH4 + + 3H2O ( 1) 4Ca0 + NO3 + 10H + 4Ca2 + + NH4 + + 3H2O 4Ca + 10HNO3 (разб) 4Ca (NO3) 2 + NH4NO3 + 3H2O Examples of some half-reactions with participation of the most typical oxidizers and reducers: Concentrated sulfuric acid SO4 2 + + 8H + 6e = S + 4H2O 75 SO42 NO3 NO3 NO3 + 4H + + 2e = H2SO3 + H2O Nitric acid + 10H + + 8e = NH4 + + 3H2O + 4H + + 3e = NO + 2H2O + 3H + + 2e = HNO2 + H2O NO3 + 2H + + e = NO2 + H2O Manganese compounds MnO4 MnO4 MnO4 Cr2O72 CrO42 + 8H + + 5e = Mn2 + + 4H2O + 2 H2O + 3e = MnO2 + 4OH + e = MnO42 Chromium compounds + 14H + + 6e = 2Cr3 + + 7H2O + 4H2O + 3e = Cr (OH) 63 + 2OH Hydrogen peroxide H2O2 + 2e = 2OH H2O2 + 2H + + 2e = 2H2O 2H + + O2 + 2e = H2O2 2H2O + O2 + 2e = H2O2 + 2OH COMPLEX COMPOUNDS Complex compounds are those which are formed of separate, rather simple chemical particles capable to independent existence (cations, anions, molecules). As examples of complex compounds we may mention Na3[Co(NO2)6]; [Cu(NH3)4]SO4; K4[Fe(CN)6]. The elementary complex compound is the ammonium cation: NH3 + H + = NH4 + 76 Structure of complex compounds One of the atoms of the complex compound, usually positively charged metal cation, occupies the central place in a complex (the central ion or a complex forming ion). Neutral molecules or ions with an opposite charge (ligands) form coordinate chemical bonds with the central ion. The central ion and ligands form inner sphere of the complex compound (a complex ion). Oppositely charged ions are electrostatically connected with the complex ion and neutralised its charge. They carry the name of outer sphere of a complex. Outer and inner spheres of a complex are connected by ionic bonds. The total number of chemical bonds formed by the central ion with the ligands, is called as coordination number of the central ion. The coordination number depends on a complexing ion charge (for monocharged ions it is usually 1, for dicharged - 4 or 6, for tricharged - 6 and above), and the size of the ion. The number of the chemical bonds formed by one ligand, is is known as its dentation. Mono, bi - and polydentate ligands are distinguished. The dentation it is defined by the number of lone electron pairs in the ligand molecule and their mutual disposition. For example, the N atom in the ammonia molecule has one lone electron pair, therefore ammonia is a monodentate ligand. The water molecule has 2 lone electron pairs, and chloride-ion has them four. However because the valence orbitals in oxygen and chlorine are in sp3-hybridization state, and are located under a corner of 10928 ’ they cannot overlap with the valence orbitals of the same central ion, therefore such ligands act as monodentate. Exceptions are polynuclear complexes, for example, Na2 [CuCl6]: Cl Na2 Cu Cu Cl Cl Cl Cl Cl 77 The nomenclature of complex compounds The name of complex compounds are similar to the names of simple salts and consists of the name of the cation and the anion. The number and names of the ligands are included into the appropriate place: [Cu (NH3) 4] Cl2 – tetra ammine copper(II) chloride; K2 [Cu (OH) 4] – potassium tetra hydroxo cupprate(II); Н [HgCl2] – hydrogen dichloromercurate(I); [Cr (NH3) 3Cl3] – trichloro triammine chromium(III). Classification of complex compounds Different types of classification of complex compounds exists. 1. Based on the charge of inner sphere: Cationic complexes: [Cr(H2O)6]Cl3, [Co(NH3)6]Cl3. Anionic complexes: K2 [HgI4], Na[Sb(OH)6]. Neutral complexes: [Pt(NH3)2Cl2]. 2. Based on the nature of the ligands: Aqua-complexes with the water molecules as ligands: [Cu(H2O)5]Cl2. Ammine complexes with the ligands as molecules of ammonia or organic ammines: [Cr(NH3)6]Cl3; [Cu(H2N-CH2-CH2-NH2)2]SO4. Hydroxo complexes: K3[Al(OH)6]. Carbonyle complexes: [Fe(CO)5]. Acido-complexes with the acidic anions as ligands, can be further devided into chloro complexes (K2[CuCl4]), cyano-complexes (K3[Fe(CN)6), nitrito-complexes (K3[Co(NO2)6), etc. 3. Based on the nature of the central ion: complexes of copper, silver, iron, chromium, etc. Dissociation of complex compounds and complex ions The majority of complex compounds are electrolytes. In solutions they are irreversibly dissociated into the inner and outer spheres (primary dsociatiation): 78 [Cu(NH3)4] SO4 [Cu(NH3)4]2 + + SO42- The complex ion is capable to the secondary dissociation as a weak electrolyte (reversible and stepwise dissociation): [Cu (NH3) 4] 2 + [Cu (NH3) 3] 2 + + NH3 2+ [Cu (NH3) 3] [Cu (NH3) 2] 2 + + NH3 2+ [Cu (NH3) 2] [Cu (NH3)] 2 + + NH3 [Cu (NH3)] 2 + Cu2 + + NH3 2+ 2+ Total process: [Cu (NH3) 4] Cu + 4NH3 Each of these reversible processes is characterised by an equilibrium constant Кinst which carries the name of an instability constant of a complex ion (stepwise or the overall): K inst [Cu 2 ][ NH 3 ] 4 2 [Cu( NH 3 ) 4 ] The less the instability constant, the mire stable is the complex ion. Instability constants are tabular values. The equilibrium constant of the reaction of formation of a complex ion at the expense of interaction of the central ion with the ligands carries the name of a formation constant or the stability constant, Kf. As reaction of formation of complex connection is opposite to its reaction диссоциации, Kf.= 1/Kinst Shifting of the equilibrium of dissociation of a complex ion submits to all rules of the le Chatelieu’s principle described previously. Reactions of complex compounds 1. Formation of complex compounds CuSO4 + 4NH3 [Cu(NH3)4]SO4 In the course of reaction, the complex ion is formed which is a weak electrolyte, therefore the reaction is irreversible. AgCl + 2NH3 [Ag(NH3)2]Cl Weak electrolytess are present both among the reactants (insoluble compound) and the products (a complex ion), therefore the reaction is 79 reversible. A direction of such reaction can be determined after calculation of the equilibrium constant of the reaction. 2. Destruction of complex compounds [Cu(NH3)4]SO4 + 4HNO3 CuSO4 + 4NH4NO3 (Complex ions are unstable in presence of strong acids and bases) [Cu(NH3)4]SO4 + H2S CuS + (NH4) 2SO4 + (NH4) 2S 3. Reactions of an exchange of the outer sphere K[Sb(OH)6] + NaCl Na[Sb(OH)6] + KCl (The reaction is possible because of formation of a precipitate) 4. Exchange reactions in the inner sphere Ligands exchange: K3 [Fe(SCN)6] + 6KCN K3[Fe(CN)6] + 6KSCN Central ion exchange: K2 [SnCl4] + CuCl2 K2 [CuCl4] + SnCl2 (Кreaction = 5103, that means the reaction is possible) Electronic structure of complex ions Interaction of lone electronic pairs of the ligands with non-filled valence orbitals of different types of the central ion leads to their hybridization. For example, the electronic structure of an ion [Cu(NH3)4] 2+ can be reflected in the following way: 2 2 6 2 6 10 1 29Cu: 1s 2s 2p 3s 3p 3d 4s 2+ 2 2 6 2 6 9 0 0 0 29Cu : 1s 2s 2p 3s 3p 3d 4s 4p 4d 3d 4s 4p 4d 4:NH3 Interaction of one s - and three p-orbitals leads to hybridization of the sp3 type (tetrahedral complex). Some valence d- and f-orbitals of the central ion may be occupied by paired or unpaired electrons. This leads to their interaction with the lone electron pairs of the ligands. This interaction is defined by the degree of penetration of the lone pairs of the ligands into the d-subshell of the central 80 ion. From this poinf of view, the ligands may be arranged in a spectrochemical series which allocate ligands of strong and weak fields: CO, CN > NO2 - > NH3 > SCN - > H2O > OH - - > F- > CL - Ligands of a strong field Ligands of a weak field Lone electron pairs of the ligands of a strong field deeply penetrate into a valence electronic shell of the central ion and cause pairing of 3delectrons. As a result, inner-orbital low spin complexes are nformed (complex compounds are formed with participation of internal 3d-orbitals, the formed complexes have no small amount of non-paired electrons): 2 2 6 2 6 7 2 27Co: 1s 2s 2p 3s 3p 3d 4s 3+ 2 2 6 2 6 6 0 0 0 27Co : 1s 2s 2p 3s 3p 3d 4s 4p 4d 3d 4s 4p 4d 4s 4p 4d [Co (NH3) 6] 3 + 3d 6:NH3 Hybridization type is d2sp3, octahedral complex. The lone electron pairs of the ligands of a weak field slightly interact with the 3d-electrons of the central ion and do not cause their pairing. Are as a result, outer-orbital high spin complexes are formed: [CoF6] 3 3d 4s 4p 4d 6:F Hybridization type is sp d , octahedral complex. 3 2 81 82 PART 2 INORGANIC CHEMISTRY 83 84 S-ELEMENTS (ALKALINE and ALKALINE EARTH METALS) Alkaline metals Li, Na, K, Rb, Cs, Fr are situated in the IA group of the periodic table. The outer spheres of these atoms consist of only one electron. So atoms of alkaline metals have a tendency to lose their valence electrons to be transformed into positively charged ions. Their oxidation number is +1. The strength of the attraction of the outer electron to the atom can be valued with the help of the ionization potential which determines the energy of removal of one electron from a neutral atom. In the IA group from Li to Fr the value of ionization potential decreases, the chemical activity of metals increases. Alkaline metals are strong reducing agents. In air alkaline metals are easily oxidized, that is why they are stored under oil. Alkaline metals are more active than hydrogen (they have negative values of redox-potentials) so they can replace hydrogen both from acids and water: 2M + HCl = 2MCl + H2 2M + 2H2O = 2MOH + H2 In aqueous solutions metal hydroxides behave as strong electrolytes and are fully dissociated: MOH M+ + OH Almost all salts of alkaline metals are water soluble. Solutions of salts containing anions of weak acids undergo hydrolysis; they are basic. Pure alkaline metals are produced by electrolysis of their melted salts. Biological significance of alkaline metals Lithium is found in the liver and lungs of animals. Large concentrations of lithium are dangerous for humans . Particles of dust and smoke containing lithium provoke malignant tumors. Sodium as NaCl is necessary for the balance of salt exchange in organisms, sodium bicarbonate is used to lower the acidity of gastric juice. In living 85 organisms potassium is situated in liver, spleen, it regulates the function of muscle cells and nervous systems. Rubidium is found in the leaves of plants (beetroot, sugar-cane, tabacoo, tea, coffee, cocoa). In animal organisms it is localized in muscles performing large load - heart muscle and pectoral muscles of birds. Cesium is found in mineral water, plants and living organisms. Compounds of rubidium and cesium are necessary for the growth of plants. Alkaline earth metals Ca, Sr, Ba, Ra are situated in the IIA group of the periodic table. Atoms of these elements have two valence electrons. While losing them, atoms of alkaline earth elements transfer into positively charged ions and gain the oxidation number +2. The presence of non-filled sublevels (d- and f-) makes Ca, Sr, Ba and Ra more chemically active and having physical properties rather different than those of Be and Mg. Metals of IIA group are less active than alkaline metals. The growth of the atomic radius, lowering of the ionization potential and of the electronegativity define the growth of their chemical activity with increase of charges of the nuclei. In the air alkaline earth metals easily form oxides. They replace hydrogen both from water and from acids: M +2HCl = MCl2 + H2 M + 2H2O = M(OH)2 + H2 The basic character of oxides and hydroxides increases with the increase of atomic radii from calcium to radium. Alkaline earth metals form different sparingly soluble salts: carbonates, phosphates, chromates, sulfates. The solubility of sulfates lowers from calcium to radium. Natural water containing soluble salts of calcium and magnum is called hard water. The presence of hydrocarbonates of calcium and magnum stipulates the temporary hardness of water. Chlorides and sulfates of these elements cause the constant hardness. The sum of temporary and constant hardness gives the overall hardness of water. 86 Biological and agricultural properties of elements of IIA group Beryllium and its compounds are toxic. The beryllium-poisoning may cause the death. Magnum compounds can be found in algae, fungi, ferns, in the tissues of animals. Magnum is a complexing ion in chlorophyll. Calcium is a constituent of the bones of vertebrates, predominantly as ortophosphate Ca3(PO4)2. Egg-shells, tests of sea animals, shells mainly consist of calcium carbonate. Organic salts of calcium play a significant role in metabolism of plants. The deficiency in calcium leads to stopping of their growth, development of rhizomes, the leaves cover by brown spots and die off. The animals suffer from rachitis, the heart activity decreases, the blood coagulability becomes worse. Calcium ions enter human organisms with milk and meat meals while magnum ions - with vegetable meal. Strontium compounds in human organisms are mainly concentrated in bones, the excess (more than 10-3 %) leads to their fragile. There are approximately 100 times less barium compounds in human bodies than those of strontium. In very small quantities barium compounds stimulate activity of marrow. In large quantities they are very toxic and provoke weakness, gastric-intestinal diseases, brain disorders. Barium chloride and carbonate are used in the agriculture as chemical weed-killers and pest-killers. ELEMENTS OF IIIA AND IVA GROUPS (P - ELEMENTS) Boron and aluminum are elements of the IIIA group of the periodic table. Atomic radius of boron is 0.91Å and the one of aluminum is 1.43Å. This great difference affects chemical properties of these elements. Ionization potential of boron is greater than that of aluminum. 87 Polarity of B-O chemical bond is small, so in solutions boron exists as BO2 and BO33 ions (acidic properties). Al-O chemical bonds have a more polar character, so in solutions aluminum exists both as Al 3+ and AlO2 ions (amphoteric properties). Salts of boric acid H3BO3 are metaborates (Ba(BO2)2) and tetraborates (Na2B4O7 - borax). Biological properties of boron and aluminum Physiological activity of boron is rather high. Together with Mn, Cu, Zn and Mo it is among five most important microelements. It concentrates in bones, teeth, muscles, marrow, liver and thyroid gland, can be found in adipose tissues of some animals, in milk and yolk of eggs. Boron inhibits the action of amilaze and proteinaze, vitamins B2 and B12, reinforces the action of insuline. Boric acid and borax are used in medicine as anticeptics. Some compounds of aluminum are also used in medicine: KAl(SO4)2 as astringent; AlOH(CH3COO)2 for desinfection; Al2(SO4)3 as coagulant. Carbon and silicon are elements of IVA group of the periodic table of elements. Their highest oxidation number is +4. At a room temperature carbon and silicon are inert elements, their activities increase with heating. At high temperatures they react with the majority of non-metals and metals. Concentrated nitric and sulfuric acids oxidize carbon into CO2, silicon can be oxidized by mixture of HNO3 and HF. Silicon can also be dissolved in alkalis: Si + 2NaOH Na2SiO3 + 2H2 Carbon (II) oxide CO is a non-salt forming oxide. Carbon dioxide CO2 has acidic properties and reversibly dissolves in water to form a weak carbonic acid: 88 CO2 + H2O H2CO3 H+ + HCO3 Salts of carbonic acid are carbonates (Na2CO3) and hydrocarbonates (NaHCO3). In aqueous solutions carbonates and hydrocarbonates undergo hydrolysis: Na2CO3 + H2O NaHCO3 + NaOH Silicic acid is even weaker than carbonic acid, the reactions of hydrolysis of its salts lead to the formation of polyanions: 2Na2SiO3 + H2O Na2Si2O5 + 2NaOH ELEMENTS OF VA AND VIA GROUPS (P-ELEMENTS) Nitrogen and phosphor are elements of VA group of the periodic table. They have 5 electrons in the outer shell (valence electrons). They may lose electrons and form positively charged ions (oxidation numbers from +1 to +5) or gain electrons and form negatively charged ions (oxidation number -3). Hydrogen compounds of nitrogen and phosphor are ammonia NH3 and phosphine (PH3). The presence of a lone electron pair of nitrogen and phosphor leads to the possibility of formation of a coordinate bond with a proton. In aqueous solutions ammonia interacts with water molecules to form ammonium hydroxide which possesses weak basic properties: NH3 + H2O NH4OH NH4+ + OH Ammonium ion has almost same properties as metallic ions, for example it can form salts. Ammonium salts decompose while heating: NH4Cl NH3 + HCl In oxides nitrogen has various oxidation states from +1 to +5. Nitric acid HNO3 is one of the strongest acids with a high oxidizing strength. Depending on the nature of a reducing agent and the concentration 89 of the acid, the NO3 group can gain from 1 to 8 electrons and transfer into NO2, NO, N2O, N2 or NH4+. Salts of nitric acid (nitrates) are water soluble. Nitrous acid HNO2 is a weak acid with redox duality. It exists only in diluted solutions and decomposes at high concentrations: 3HNO2 HNO3 + 2NO + H2O In contrast with nitric acid, phosphoric acid has no oxidizing properties. Phosphates form soluble complexes with a lot of metal ions. Sulfur is situated in the VIA group of the periodic table and has 6 valence electrons. Its oxidation numbers are +4 , +6 and -2. Sulfuric acid H2SO4 (conc.) is a strong oxidizing agent and hydrogen sulfide and its salts (sulfides) are good reducing agents. Sulfurous acid H2SO3 and its salts (sulfites) have redox duality. The strength of acids containing sulfur increases with the increase in oxidation number of sulfur. Biological importance of sulfur Sulfur is a composite of most important aminoacids. Metal sulfates are used in medicine: CaSO4 as a stuff for plasters, BaSO4 in rontgenoscopy of stomach, MgSO410H2O as purgative. Some antibiotics have sulfur compounds as composites. ELEMENTS OF THE VIIA GROUP (HALOGENS) Atoms of halogens have 7 electrons in the outer shell (ns2np5) which determine their chemical activity. Halogens are strong oxidizing agents. Their activities increase with decrease in ionic radii: fluorine is the strongest oxidizing agent. 90 The properties of fluorine differ from those of other halogens. As its atoms have no empty d-orbitals in the outer sphere, it can’t exist in excited states, and its only oxidation number is -1. Atoms of chlorine, bromine and iodine have a vacant d-orbital in their outer spheres, so 3 electrons may be unpaired to form 3 excited states. Possible oxidation numbers of these elements are -1, +1, +3, +5, +7. Bond energies in the molecules of halogens increase with decrease in atomic numbers: ECl-Cl > EBr-Br > EI-I. ECl-Cl = 57.8 kcal/mol EBr-Br = 46.1 kcal/mol EI-I = 36.2 kcal/mol The molecule of fluorine has the minimal bond energy which can be explained by the features of electronic configuration of fluorine comparing with other halogens. Hydrogen compounds of halogens are colorless gases. The bond energies decrease from HF to HI. Their aqueous solutions possess acidic properties, HI is the strongest one among them. Reducing properties of halogen hydrides increase with increase of charge of nuclei. Oxygen containing compounds of halogens are strong oxidizing agents. Interactions of halogens with water can be expressed as: F2 + H2O 2HF + [O] [O] + F2 F2O X2 + H2O HX + HOX (X = Cl, Br, I) The equilibria of these interactions are shifted to the left. In alkaline solutions the reactions become irreversible: Cl2 + 2NaOH NaCl + NaClO + H2O 91 TRANSITIONAL ELEMENTS (d-ELEMENTS) d-Elements are situated in B-subgroups of the periodic table. Their valence electrons are those of s-sublevel of the outer shell and of the unfilled d-sublevel. The presence of 1 or 2 electrons on the outer shell of all d-elements stipulates their metallic properties. d-Electrons take part in the formation of chemical bonds, so different oxidation states are known for d-elements. High values of oxidation numbers are typical only for delements with non-paired d-electrons (first 5 elements of each transitional series). The values of a first ionization potential of d-elements of one transitional series increases with increase in charge of the nuclei. d-Elements in the same oxidation states usually have similar properties. For example, all hydroxides of M(OH)2 and M(OH)3 types are weak bases which are sparingly soluble in water. Sulfide and carbonate ions form precipitates with M3+ and M2+ ions. All d-elements are good complexing agents. Ions of d-elements in their higher oxidation states possess acidic properties and exist in solutions as anions: VO3 , CrO42 , Cr2O72 , MnO42 , MnO4 . Biological significance of d-elements All derivatives of Cr (VI) are strongly toxic and provoke ulcers and lung cancer while breathing. Manganese can be found in both animals and plants tissues. Addition of small quantities of manganese to fertilizers increase the crop capacity of some plants (maize, sugar-beet, potatoes and others). Iron is a catalyst of respiratory processes. Human bodies contain 4g of iron, about 57% as haemoglobin. Soils contain from 1 to 5% of iron compounds. Deficiency of iron provokes growing leaves pale. Biological role of cobalt in living organisms is connected with circulatory system. An antianaemic vitamin B12 contains 4.35% of cobalt. Cobalt compounds also oppress malignant tumors. 92 Plants usually contain 10 4 % of zinc. Small amounts of zinc are necessary for normal growth of animals and fruiting of plants. Mercury and its compounds are very toxic and provoke disturbance of cardiac and stomach activities and weakening of memory. ELECTROCHEMICAL CORROSION OF METALS The metals are good reducing agents. Their reducing ability is described by the redox potential (Mn+/M0). The less the potential is, the stronger are the reducing activities of the metal. In the case of a macro-contact of two metals in presence of an electrolyte (or their micro-contacts which take place in alloys), a process of electrochemical corrosion, i.e. destruction of metals in the electrolyte media, occurs. The stronger reducing metal (or a metal of a higer activity or less (Mn+/M0) value) charges negatively because of the partial transmission of the metal cations in the electrolyte medium: M0 – n e Mn+ The liberated electrons move to the surface of a less active metal and interact with the ions or molecules present in the electrolyte. Depending on the pH of the medium, two processes are possible: a) In acidic media (pH<7), the hydrogen depolarization occurs: 2H+ + 2 e H2; b) In neutral (pH = 7) or in alkaline media (pH>7), the oxygen depolarization becomes possible: O2 + 2H2O + 4 e 4OHExample. Corrosion of a Cu/Fe contact in presence of NaOH. (Cu2+/Cu0) = +0.34 V; (Fe2+/Fe0) = –0.44 V (Fe is more active than Cu). Fe – 2 e Fe2+ (x2) Cu surface: O2 + 2H2O + 4 e 4OH - (x1) 2Fe + O2 + 2H2O 2Fe(OH)2 93 94 PART 3 ORGANIC CHEMISTRY 95 96 THE THEORY OF CHEMICAL STRUCTURE OF ORGANIC COMPOUNDS Organic compounds are carbon compounds (other than the simplest) in which it exhibits a valence of IV. Organic chemistry - a chemistry of hydrocarbons and their derivatives. The carbon atom in organic compounds is in an excited state, and has four unpaired electrons : 2 2 2 2 1 3 → 6С 1s 2s 2p 6С* 1s 2s 2p sp3 The carbon atom in the excited state is capable: 1) to form strong bonds with other atoms, which leads to the formation of chains and cycles ; 2) due to the different types of orbital hybridization to form simple, double and triple bonds between carbon and other atoms (H, O, N, S, P et al.) 3) connect to four different atoms, which leads to the formation of branched carbon chains . Types of hybridization of carbon atoms in organic compounds sp3 - hybridization All four valence orbitals are involved in the hybridization. Valence angle 109o28 '(tetrahedron). Carbon atoms form only simple (σ) bonds saturated compounds. sp2 - hybridization Three hybrid and one non-hybrid orbitals are formed. Bond angle 120° (flat structure, an equilateral triangle) . Hybrid orbitals form σ- bonds and a non-hybrid orbital forms a π-bond. 97 sp - hybridization Formation of two hybrid and two non-hybrid orbitals. Bond angle of 180o (linear structure). The carbon atom in sp- hybridization takes part in the formation of two double bonds or one triple bond. The theory of the structure of organic compounds was formulated by AM Butlerov in 1861 and includes the following provisions: a) All the atoms in the molecule are linked together in a strict sequence according to their valences. The chemical structure determines the way of connecting the atoms in a molecule. b) Properties of organic compounds depend not only on the qualitative and quantitative composition of substances , but also on the chemical structure of the molecule. c) Atoms in the molecule have a mutual influence on each other, i.e., properties of atomic groups in the molecule may vary depending on the nature of the other atoms in the molecule. A group of atoms that determines the chemical properties of organic molecules is called a functional group. d) Each organic compound having only one chemical formula. Knowing the chemical formula, we can predict the properties of the compound, and by studying its properties, in practice, to establish the chemical formula. organic molecule carbon chain atoms carbon skeleton 98 + Functional Group identifies the classes of organic compounds Types of the carbon skeleton : - Acyclic : • branched ; • normal (linear) . - Cyclic: • carbocyclic (cycle of only carbon atoms); • heterocyclic (except carbon atoms in the cycle includes several other atoms - nitrogen , oxygen, sulfur) . Types of carbon atoms in the hydrocarbon chain СН3 Н3С-СН2-СН-С- СН3 СН3 - СН3 Primary carbon atoms (connected in a chain with only one carbon atom is the terminal); Secondary carbon (bonded to two adjacent carbon atoms located in the middle of the chain); A tertiary carbon atom (located on the branched carbon chain connected to three carbon atoms) A tertiary carbon atom (located on the branched carbon chain connected to three carbon atoms); Functional group is a special group of atoms, which determines the chemical properties of the compounds. Examples of functional groups: -ОН – hydroxyl group (alcohols, phenols); С=О – carbonyl group (ketone, aldehyde); ОН С - A carboxyl group (carboxylic acid); О 99 -NH2 – Amino group (amine); -SH – Thiol group (thioalcohols) organic compound Composition Properties chemical structure The atoms that make up the organic compound can be connected in different ways to form molecules. For example, a compound of C2H6O can respond to two chemical compounds having different physical and chemical properties: H H | | H C C O H | | H H H H | | H C O C H | | H H ethanol (b.p.= 78C, Reacts with Na) Dimethyl ether b.p. = –24C, does not react with Na) Isomers are compounds having the same composition but different chemical structure. Isomers have different chemical properties. TYPES of ISOMERISM 1 structural isomers Isomers of the carbon chain: CH3 CH2 CH2 CH3 CH3 CH CH3 CH3 Butane 100 Methylpropane Positional isomers of multiple bonds: СН2=СН–СН2–СН3 СН3–СН=СН–СН3 Butene-1 Butene-2 Positional isomers of the functional group: CH3 CH2 CH2 OH CH3 CH CH3 OH 1-propanol Interclass isomerism: СН3–СН=СН–СН3 2-propanol HC CH HC CH butene-2 Cyclobutane 2-stereoisomers Geometric (spatial, cis-trans isomers of compounds with double bonds) H H C H3C H C CH3 C CH3 cis-butene-2 C H3C H trans-butene-2 Geometrical isomerism is possible in the event that each of the carbon atoms involved in the formation of a double bond, has a different substituents. Thus, for a butene-1 CH2=CH-CH2-CH3, geometric isomerism is not possible, since one of the carbon atoms at the double bond has two identical substituents (hydrogen atoms). Geometric (spatial, cis-trans isomers of cyclic compounds 101 H H CH3 H H H H CH3 H CH3 H CH3 cis- 1,2-dimethylcyclopropane trans 1,2-dimethylcyclopropane Geometrical isomerism is possible in the case when at least two carbon atoms forming the ring, have different substituents Optical isomerism: Optical isomerism is stereoisomerism view caused the chirality of molecules. In nature, there are compounds which correspond the two hands of one person. One of the properties of these compounds is their incompatibility with its mirror image. This property is called chirality (from the Greek. "Sheir" the hand). The optical activity of molecules detected when exposed to polarized light. If the solution through an optically active substance skip polarized beam light, it will rotate the plane of polarization. Optical isomers are designated using prefixes d- and l-. For ease of reading optical isomerism Fischer proposed flat formula reflecting the configuration of the chiral centers: CHO OH CH2OH D- glyceraldehyde Mirrow H CHO HO H CH2OH L- glyceraldehyde The optical isomerism is possible, if the molecule contains one or more carbon atoms in sp3-hybridization condition that contain four different substituents. 102 CONFORMATIONAL (SWING) ISOMERISM As a result of the tetrahedral configuration of chemical bonds in the hydrocarbon chains of saturated compounds are non-linear, and are arranged in a zigzag pattern space; angle between the bonds corresponds to the tetrahedral: Free rotation occurs around a simple C-C bond, so the chain in space can take many forms or conformations. Free rotation inhibited the interaction of the hydrogen atoms on adjacent carbon atoms: H HH H H H H H H eclipsed conformation H H H inhibited conformation Conformational isomerism cycloalkanes linked with a nonplanar structure of major cycles (containing more than 6 carbon atoms) conformation "tub" conformation "chair" Homologous series The special features of organic compounds can also be considered as the existence of a homologous series, in which each successive term can 103 be generated from the previous addition of one specific for a given number of groups of atoms. For example, in the homologous series of saturated hydrocarbons such group is CH2. Homologues have similar chemical properties, while the physical properties change somewhat with increasing molar mass compounds: СН4 Methane С2Н6 ethane, С3Н8 propane, С4Н10 butane, С5Н12 pentane Types of Reactions in Organic Chemistry Most organic compounds are characterized by a relatively low rate of chemical reactions under normal conditions . This is due to the high strength of the covalent bonds of carbon - carbon bond; Carbon - hydrogen; carbon - oxygen and others. Among the values of electronegativity carbon occupies an intermediate position between the typical oxidants and reducing agents, so that the difference in electronegativity and the polarity of carbon bonds with other atoms is small (the chemical bonds in organic compounds of low polarity and do not dissociate in solution , that is, the majority of organic compounds - non-electrolytes) . Types of fracture (splitting) of chemical bonds : Homolytic decomposition of 1 - equivalent splitting the total electron pair to form radicals : СН4 → •СН3 + •Н 104 Stabilization of the formed radical is due to interaction with neighboring carbon bonds. Radicals formed by the tertiary carbon atom stable rest. Heterolytic decay - unequal splitting of the total electron pair to form carbocations and carbanions : СН3Cl → СН3+ + Cl– carbocation СН3MgCl → СН3– + MgCl+ carboanion Types of reactions in organic chemistry: 1 Substitution of (S). 2 Addition (A). 3 Splitting (elimination) (E). 4 Rearrangement (isomerization). Attacking particles: • radical (R). • Negative and positive particles: - Nucleophile (N) (-); - An electrophile (E) (+). HYDROCARBONS Hydrocarbons are organic compounds that contain only carbon and hydrogen atoms. Their characteristic feature is the lack of functional groups. Properties of hydrocarbons determined by the structure of the hydrocarbon radical. Saturated hydrocarbons: 105 ALKANES Alkanes (CnH2n+2, acyclic, all carbon atoms are in sp3-hybridization condition). Chemical properties. Low polar covalent bonds of alkanes and very durable, so saturated hydrocarbons in chemical reactions are inactive. The main types of chemical reactions of alkanes are: 1. Halogenation (radical substitution reaction, SR): 2 CH3 CH2 CH3 + Cl2 h CH3 CHCl CH3 + CH3 CH2 CH2Cl + 2HCl The formed alkyl halides are used for producing alkanes with a longer carbon chain, as well as compounds of other classes (alkenes, alcohols) CH3 CH3 CH Br + 2 Na + Br CH CH3 CH3 CH3 to 2 NaBr + H3C CH CH CH3 CH3 (Wurtz reaction) H2O H3C CH CH3 + NaBr OH H3C CH CH3 + NaOH Br o t C2H5OH H3C CH CH2 + NaBr + H2O 2. Sulfochlorination. RН + SO2 + Cl2 → HCl + RSO2Cl RSO2Cl + NaOH → NaCl + RSO3Na (sodium alkyl sulfonate, based on synthetic detergents) 3. Nitration (Konovalov’s reaction): 106 t, P CH3 CH2 CH3 + HNO3 CH3 CH CH3 + H2O NO2 2-nitropropane 3. Oxidation СnH2n+2 + (1,5n+0,5)O2 nCO2 + (n+1)H2O (burning); 4. Isomerization CH3 CH2 CH2 CH3 o 100 C , AlCl3 CH3 CH CH3 CH3 5. Dehydrogenation: CH3 CH2 CH2 CH3 o 300-450 C CH3 CH CH CH3 + H2 Al2O3 или Cr2O3 CH3 CH2 CH2 CH3 o 400-600 C Al2O3 CH2 CH CH2 CH3 + H2 CH2 CH CH CH2 + 2 H2 CH3 CH2 CH2 CH2 CH2 CH3 o 300 C Pt + 4H2 CH3 CH2 CH2 CH2 CH2 CH2 CH3 o 300 C Pt CH3 + 4H2 6. Cracking: СH4 1000oC C + 2H2; С18H38 2CH4 t>1500oC C2H2 + 3H2; t C9H18 + C9H20 107 CYCLOALKANES Cycloalkanes (CnH2n are cyclic, all the carbon atoms are in sp3hybridization). The angles between the bonds in small rings - C3H6 (600) and C4H8 0 (90 ) are less than required by the geometry of the carbon atom in the state of sp3-hybridization (109.50), so they are unstable, and characterized by the addition reaction, coming from breaking up the cycle: С4H8 + H2 С3H6 + HBr Pd, t C4H10; t C3H7Br; С3H6 + Br2 t C3H6Br2. Cycloalkanes based on 5, 6 and 7 carbon atoms (normal cycle) are chemically inactive, same as alkanes, they are characterized typical substitution reaction С6H12 + Сl2 h C6H11Cl + HCl. Unsaturated hydrocarbons ALKENES Alkenes (CnH2n are acyclic compounds with thecarbon atoms in the states of sp3- and sp2 hybridization with one double bond). Addition reactions (in electrophilic mechanism, AE) 1. Halogenation: CH2=CH2 + Br2 CH2Br–CH2Br (discoloration of bromine water - an aqueous solution of bromine is a qualitative reaction of compounds with double and triple bonds). When the temperature rises above 500 ° C in the halogenation of alkenes are substitution reaction 108 СН3-СН=СН2 + Br2 540C СН3-СBr=СН2 + HBr 2. Hydrogenation and dehydrogenation: CH2=CH2 + H2 CH2=CH–СН2–СН3 , Pt , Pd ил и Ni t CH3–CH3 , Cr2O3 t CH2=CH–СН=СН2 3. Hydrohalogenation: СH3–CH=CH2 + HBr CH3–CHBr–CH3 (the reaction proceeds via Markovnikov's rule: hydrogen is added to the double bond to the most hydrogenated carbon atom). In the presence of hydrogen peroxide or other peroxide compounds reaction mechanism is changed to a radical reaction and gidrogalogenirovaniya goes against the rules of Markovnikov (Kharasch effect): Н2О2 СН3-СН=СН2 + HCl СН3-СН2-СН2Cl (АR) 4. Hydration: СH3–CH=CH2+H2О 3 PO4 H CH3–CH(ОН)–CH3 Oxidation reactions СnH2n + 1,5nO2 nCO2 + nH2O (burning); 3CH2=CH2 + 2KMnO4 + 4H2O 3HO–CH2–CH2–OH + 2MnO2 + 2KOH (soft oxidation in a neutral medium); 5СH3–CH=CH–CH3 + 8KMnO4 + 12H2SO4 10CH3COOH + 8MnSO4 + 4K2SO4 + 12H2O; 5 CH 3 C CH 2 + 8KMnO4 + 12H2SO4 | CH3 109 5 CH 3 C O + 5CO2 + 8MnSO4 + 4K2SO4 + 17H2O | CH3 (hard oxidation in acidic medium). Polymerization n(CH3–СH=CH2) p, t, катализатор CH CH 2 | CH 3 n Polymer is a macromolecule consisting of a large number of recurring units called monomers. The number of repetitions of the monomers in the chain (n) is called the "degree of polymerization"). In the polymerization process, a mixture of macromolecules with different degrees of polymerization is usually obtained, so the polymers are not characterized by a fixed melting point, and melt in the temperature range. The polymerization reactions associated with rupture of the double bonds in the molecules of monomers. ALKADIENES Alkadienes (CnH2n-2 are acyclic. Their carbon atoms are in the sp2 and sp -hybridization states with two double bonds). Only conjugated dienes (1,3-dienes) are of biological significance 3 110 Dienes having various substituents on the carbon atoms at the double bond, like alkenes exhibit cis-trans isomerism. The chemical properties of conjugated alkadienes Electrophilic addition reaction and the polymerization of conjugated alkadiene occur, usually at positions 1 and 4: CH2=CH–CH=CH2 + Br2 CH2Br–CH=CH–CH2Br CH2Br–CH=CH–CH2Br + Br2 CH2Br–CHBr–CHBr–CH2Br CH2=CH–CH=CH2 + 2Br2 CH2Br–CHBr–CHBr–CH2Br; CH2=CH–CH=CH2 + HBr CH3–CH=CH–CH2Br CH3–CH=CH–CH2Br + HBr CH3–CHBr–CH2-–CH2Br CH2=CH–CH=CH2 + 2HBr CH3–CHBr–CH2-–CH2Br. Hydration is not characteristic of conjugated alkadienes, as by heating in presence of acids (condition for hydration of alkenes) their polymerization occurs, which leads to the formation of rubbers: p, t, катализатор nCH2=CH–CH=CH2 (–CH2–CH=CH–CH2–)n ALKYNES Alkynes (cnH2n-2 acyclic carbon atoms are in the states sp and sp3hybridization within the molecule contains one triple bond). Addition reactions Ni , t Hydrogenation: СHCH + 2H2 CH3–CH3 Halogenation CHCH + Br2 CHBr=CHBr; 111 CHBr=CHBr + Br2 CHBr2–CHBr2 or CHCH + 2Br2 CHBr2–CHBr2 Hydrohalogenation: CHCH + НBr CH2=CHBr; CH3–ССН + HBr СH3–CBr=CH2; CH3–CBr=CH2 + HBr CH3–CBr2–CH3 Hydration (Kucherov reaction): HC CH + H2O HgCl2 O [H2C CH] H3C C H OH HC C CH3 + H2O HgCl2 [H2C CH CH3] H3C OH C CH3 O Polymerization СН2=СН–CCH 2СНСН CuCl , NH 4Cl 3CH CH o 600 C, C активир. (benzene); Acidic properties of alkynes Terminal alkynes have acidic properties - the substitution of the hydrogen atoms have sp-hybridized carbon atom of the metal: СНСН + 2[Ag(NH3)2]OH Ag–CC–Ag + 4NH3 + 2H2O; CH3–CCH + CuCl + NH4OH CH3–CC–Cu + 2NH3 + H2O; CH3–CC–CH3 + Ag2O (NH3) CH3–CCH + NaH CH3–CC–Na + H2; CH3–CC–Na + H2O CH3–CCH + NaOH CH3–CC–Na + СH3–CH2I NaI + CH3–CC–CH2–CH3 112 AROMATIC HYDROCARBONS Aromatic hydrocarbons or arenes (CnH2n-6 are cycliccompounds with the carbon atoms in sp2-hybridization state, a molecule contains conjugated system of double bonds). H HC C C H C C H C H H или The representatives of the homologic series of arene are: С6Н6 С7Н8 CH3 benzene, methylbenzene (toluene) С8Н10 CH2 CH3 CH3 ethylbenzene CH3 CH3 CH3 H3C CH3 1,21,31,4dimethylbenzene dimethylbenzene dimethylbenzene (o-xylene) (meta-xylene) (p-xylene) CH CH2 vinylbenzene (styrene) H3C CH CH3 isopropylbenzene (cumene) 113 In condensed aromatics two adjacent "spliced" cycle have two atoms in common. Thus, three types of coupling loops. naphthalene anthracene phenanthrene Chrysene pyrene Benzpyrene Chemical properties . Conjugated π- electron system of benzene and its homologues is an energetically favorable state, so its destruction occurs with great difficulty. Addition reaction to arenes are not typical, more common are electrophilic substitution reactions (SE). They occur in presence of catalysts : Fe, trivalent metal salts - Lewis acids. Lewis acids are particles that are capable of accepting lone electron pairs: AlCl3 + Cl– → [AlCl4]– Chemical properties of benzene 1. Halogenation: 114 + Br2 Br AlBr3 или FeBr3 + HBr 2. Nitration NO2 H2SO4 (конц.) + HNO3 (HO NO2) + H2O . 3.Alkylation CH3 AlCl3 + CH3Cl + HCl ; + CH2 CH CH3 CH CH3 H3PO4 CH3 ; CH CH3 H2SO4 + CH3 CH CH3 CH3 + H2O OH . 4. Addition reactions: + 3 H2 P, t, Ni (cyclohexane); H + 3 Cl2 h Cl H Cl H Cl Cl H Cl H Cl H (hexachlorocyclohexane). 115 5. Oxidation: 12С6Н6 + 15О2 12СО2 + 6Н2О (burning). Benzene can not be oxidized by potassium permanganate. Chemical properties of benzene homologues The alkyl substituents on the benzene ring slightly increase the activity of the benzene ring and activy the substitution at positions 2, 4, and 6 relative to the alkyl substituent (ortho and para-substitution) CH3 Br CH3 + HBr + Br2 AlBr3 или FeBr3 CH3 + HBr Br Nitration of toluene leads to the formation of the trinitro-substituent: CH3 CH3 H2SO4 O2N NO2 + 3 H2O + 3 HNO3 NO2 (trinitrotoluene). When irradiated with the UV light, the benzene homologues undergo the reaction of substitution in the side chain, with the greatest activity are carbon atoms directly connected to the benzene ring: CH2 CH2 CH3 + Br2 CH CH2 CH3 h Br + HBr . 116 Benzene homologues burn to form carbon dioxide and water, and can be oxidized with potassium permanganate in the side chain to form a benzoic acid or its salts: С6Н5СН3 + 9О2 7СО2 + 4Н2О (burning); 5С6Н5CH3 + 6KMnO4 + 9H2SO4 5C6H5COOH + 6MnSO4 + 3K2SO4 + 14H2O; t С6Н5CH3 + 2KMnO4 C6H5COOK + 2MnO2 + KOH + H2O; 5С6Н5CH2CH3 + 12KMnO4 + 18H2SO4 5C6H5COOH + 5CO2 + 12MnSO4 + 6K2SO4 + 28H2O; t С6Н5CH2CH3 + 4KMnO4 C6H5COOK + K2CO3 + 4MnO2 + KOH + 2H2O. Characteristics of the chemical properties of condensed rings The mutual influence of the two condensed rings (e.g., in naphthalene) leads to increased activity of the hydrogen atoms in the α-position: Br Br2 HBr 117 KMnO4 COOH COOH ORGANIC COMPOUNDS WITH FUNCTIONAL GROUPS Functional group is a group, which defines the affiliation to a class of organic compounds and chemical properties. Depending on the number and nature of functional groups, the following types can be distinguished: Monofunctional compounds which have only one functional group of alcohols, aldehydes, ketones, carboxylic acids, ethers, esters, phenols, amines, etc. Multifunctional compounds that have multiple identical functional groups (polyhydric alcohols, polybasic acids and others.). Polyfunctional compounds that contain several different types of functional groups: amino acids, carbohydrates, etc. AMINES Amines are organic derivatives of ammonia in which one or more hydrogen atoms are replaced by a hydrocarbon radical: 118 H3C CH3–NH2 CH3–NH–CH2–CH3 methylamine (primary amine) methylethylamine (secondary amine) N CH3 H3C trimethylamine (tertiary amine) In the aromatic amines, the nitrogen atom is connected to the benzene ring: NH2 H CH3 N H 3C CH3 N N Phenylamine (Aniline) Methylphenylamine Dimethylphenylaminee Triphenylamine (N-methylaniline) (N, Ndimethylaniline) Chemical properties 1: Basic properties. The nitrogen atom in the ammonia and amines has an lone electron pair able to bind hydrogen ions due to the donoracceptor interactions. CH3–NH2 + Н2О CH3–NH3+ ОН– methylamine methylammonium hydroxide CH3–NH2 + НСl CH3–NH3+ Cl– Saturated amines are stronger bases than ammonia; secondary amines are stronger than the primary ones. Amine salts are decomposed by the action of strong inorganic bases: CH3–NH3Cl + NaOH CH3–NH2 + NaСl + H2O. 119 Aromatic amines are much weaker than ammonia and aliphatic amines. This is due to the fact that the lone electron pair of the nitrogen atom is partially shifted to the aromatic ring by conjugation with the πelectrons of the benzene rings and are therefore less available for interaction with the proton: С6H5NH2 + H2O ; C6H5NH2 + HCl C6H5NH3+ Cl– aniline phenylammonium chloride 2. Reactions of burning. Combustion of amines is not accompanied with the formation of nitrogen oxides but nitrogen as a simple substance: 4CH3NH2 + 9O2 4CO2 + 10H2O + 2N2 or CnH2n+3N + (1,5n+0,75)O2 nCO2 + (n+1,5)H2O + 0,5N2. 3. Interaction of the amino group of aromatic amines with the benzene ring. The amino group facilitates the electrophilic substitution reaction of hydrogen atoms on the benzene ring and orients them in the ortho and para positions. Aniline discolours the bromine water. At the same time, a white precipitate isd formed. This is a qualitative reaction for aromatic amines: NH2 NH2 Br Br + 3Br2 + 3HBr Br ALCOHOLS Alcohols are organic compounds containing one or more hydroxyl groups (OH) attached to a hydrocarbon radical. 120 Monohydric saturated alcohols Monohydric alcohols contain one hydroxyl group bonded to a saturated hydrocarbon radical. The general formula : R-OH, CnH2n+1OH. For alcohols, regioisomers of the functional group and isomers of the carbon skeleton are known. Alcohols in which the hydroxyl group is in the 1st position (attached to a primary carbon atom), are called the primary alcohols. If the hydroxyl group is located in the middle of the chain (attached to a secondary carbon atom), it is called a secondary alcohol. Joining hydroxyl group to a tertiary carbon atom results in the formation of tertiary alcohols. primary secondary tertiary alcohols While naming alcohols, the position of the OH group and the end “ol” are used: CH3 OH CH3 CH2 OH CH3 CH2 CH2 OH methanol ethanol 1-propanol CH3 CH CH3 CH3 CH2 CH2 CH2 OH OH 2-propanol 1-butanol CH3 CH CH2 OH CH3 1-methylpropanol 121 CH3 CH CH2 CH3 CH3 CH3 OH C CH3 OH 2-butanol 2-methylpropanol Chemical properties. Characteristics of the chemical properties of alcohols are defined by the presence of a functional group OH. Electronegative oxygen atom pulls on shared electron pairs, thus acquiring a negative charge, and the neighboring carbon and hydrogen atoms produce positive charges. Displacement of the electron pairs leads to the formation of polar bonds which are less stable and more reactive . Reactions involving the cleavage of the RО–Н bond а). 2С2H5OH + 2Na H2+ 2C2H5ONa sodium ethoxide С2H5ONa + Н2O С2H5OH + NaOH (hydrolysis of alkoxides); С2H5OH + NaОН (monohydric alcohols do not react with alkalis). b). Formation of esters: С2H5OH + СН3СООН H+, t Н2О + СН3СООС2Н5 (ethyl acetate); С2H5OH + HNO3 (HO–NO2) Н2О + С2Н5ОNO2. 122 Reactions involving the cleavage of the R–ОН bond С2H5OH + HCl C2H5Cl + H2O; а). b). Intramolecular dehydration: t 180o C , H 2 SO4 СH3–СН2–OH H2O + СН2=СН2 180o C , H 2 SO4 СH3–СН2–СН(OH)–СН3 t СH3–СН=СН–СH3 + H2O (reaction proceeds according to the Zaitsev’s rule: the H atom is detouched from the least hydrogenated carbon atom). c). Intermolecular dehydration (formation of ethers) CH3 CH2 OH t 140o C , H 2 SO4 H2O + CH3 CH2 O CH2 CH3 CH3 CH2 OH (dimethyl ether). Ethers form a separate class of organic compounds isomeric to saturated monohydric alcohols (interclass isomerism). Oxidation of monohydric alcohols a). burning СnH2n+2O + (1,5n +0,5)O2 nCO2 + (n +1)H2O. b). Air oxidation in the presence of catalysts: CH3 CH2 OH 300 500 o C , Cu CH3 O C + H2 H (primary alcohol) (aldehyde) 500 C , Cu 300 o CH3 CH CH3 OH CH3 C CH3 + H2 O 123 (secondary alcohol) (ketone) c). Oxidation by copper oxide (II): CH3 CH2 OH + CuO t CH3 O C H (primary alcohol) + H2O (aldehyde) d). Oxidation with potassium permanganate or potassium dichromate at room temperature (primary alcohols are oxidized to carboxylic acids, secondary alcohols - to ketones, tertiary alcohols are not oxidized under these conditions) 5СH3OH+6KMnO4+9H2SO4 5CO2+6MnSO4+3K2SO4+19H2O; 5СH3–СН(OH)–СН3+2KMnO4+3H2SO4 5СH3–СO–СН3+2MnSO4+K2SO4+8H2O. Monohydric unsaturated alcohols HC C CH2 1-propenol (allyl alcohol) OH 1-propynol (propargyl alcohol) 3-phenyl-1-propenol (cinnamic alcohol) Structural isomers of unsaturated alcohols are related to the structure of the carbon chain and the positions of multiple bond and a hydroxyl group. The multiple bond and OH groups can not be attached to the same C atom because of the isomerization into aldehydes or ketones: O CH2 CH OH CH3 C H vinyl alcohol 124 ethanal . Unsaturated alcohols possesst the properties of both alcohols and unsaturated compounds. Polyhydric alcohols (polyols) Polyhydric alcohols contain several hydroxyl groups attached to different carbon atoms. Attaching multiple hydroxyl groups to the same carbon atom is impossible, since the dehydration process occurs: O OH CH3 CH OH CH3 OH C OH OH H2O + CH3 C H O H2O + CH3 C OH Examples of polyhydric alcohols: H2C CH2 HO OH ethanediol (ethylene glycol) CH2 CH CH2 OH OH OH diatomic propantriol (glycerol) CH2 * * CH CH CH CH2 CH2 * * * * CH CH CH CH CH2 OH OH OH OH OH OH OH OH OH OH OH xylitol sorbitol Chemical properties of polyols 1. Acidic properties CH2 OH + 2Na CH2 OH CH2 ONa CH2 ONa + H2 ; 125 CH2 OH CH2 OH CH2 ONa + NaOH CH2 OH 2 Qualitative reaction to polyols: CH2 OH _ CH2 OH OH CH O 2 CH OH + Cu(OH)2 Cu CH2 OH CH2 O H blue precipitate + H2O . (glycolate) H O CH2 + 2H O 2 O CH HO CH2 dark blue solution (Cu glycerate) 3 Formation of full and partial esters with inorganic or organic acids: CH2 OH CH OH CH2 O CO CH3 + 3CH3COOH CH2 OH CH2 O CO CH3 CH2 OH CH OH CH2 OH CH2 OH + 3HNO3 (HO NO2) + H3PO4 (HO PO3H2) 4. Dehydration of polyols 126 ; CH2 O NO2 CH2 OH CH OH CH O CO CH3 + 3H2O CH O NO2 + 3H2O CH2 O NO2 (nitroglycerin) CH2 OH CH OH CH2 O + H2O PO3H2 . (dioxane) PHENOLS Phenols are aromatic organic compounds in which the hydroxyl groups of the molecules are linked to carbon atoms of the aromatic ring: OH OH OH OH CH3 CH3 phenol ortho-cresol meta-cresol CH2OH benzyl alcohol (does not possess phenolic properties) OH CH3 para-cresol OH α-naphthol OH β-naphthol OH OH OH OH OH catechol resorcinol hydroquinone 127 OH OH OH HO OH OH HO OH OH pyrogallol phloroglucinol oxyhydroquinone Chemical properties of phenols 1. Reactions of the OH group a) Acidic properties 2C6H5OH + 2Na H2 + 2C6H5ONa (sodium phenolate); C6H5OH + NaOH C6H5ONa + H2O; C6H5ONa + H2O + CO2 C6H5OH + NaHCO3 (acidic properties to phenol are less than that of carbonic acid); ONa OH OH + CO2 COONa H+/H O 2 125oC COOH 6 àòì ô åí î ëÿò í àòðèÿ ñàëèöèëî âàÿ êèñëî òà ñàëèöèëàò í àòðèÿ sodium phenolate sodium salycilate salycilic acid _ e 2 C6H5OH + FeCl3 + 4 H2O _ +3 [Fe(H2O)4 C6H5OH C6H5O]Cl2 + HCl +2 [Fe(H2O)4 C6H5OH C6H5O]Cl2 ô èî ëåòî âû é Violet staining solutions in the presence of iron chloride (III) is a qualitative reaction on phenols. 128 b) Formation of ethers and esters С6Н5ОН + СН3СООН O C6H5ONa + CH3 C Cl O CH3 C O C6H5 + NaCl C6H5ONa + R–Br C6H5OR + NaBr c) Oxidation OH O OH [O] O [O] O OH O benzoquinone 1. Reactions of the benzene ring (similar to the aromatic amines): Br Br + r2 + 3 HBr Br O2N NO2 + HNO3 + 3 H2O NO2 129 Trinitrophenol (picric acid) is a yellow crystalline solid, the acidic strength approaches inorganic acids). Polycondensation with formaldehyde (resins formation): OH OH O кислотные или щелочные катализаторы +nH C H (n+1) CH2 OH CH2 + n H2O n-1 ALDEHYDES AND KETONES Aldehydes and ketones are carbonyl containing compounds: O C . O H C H methanal (formaldehyde or formic aldehyde) C CH2 CH2 ethanal (acetaldehyde or acetaldehyde) C H butanal (butyraldehyde) CH3 C CH3 O propanone (acetone) CH2 C CH C H O CH2 methylpropanal (Isobutyraldehyde) CH3 CH2 C CH3 O butanone H propanal (propionaldehyde) O CH3 CH3 130 CH3 H O CH3 O O CH3 CH C H propenal (acrolein) CH3 CH2 CH2 C CH3 O 2-pentanone CH3 CH2 C CH2 CH3 CH3 C C6H5 CH3 O CH3 CH C CH3 O O 3-pentanone C6H5 C methylbutanone O methylphenylketone (acetophenone) C6H5 C C6H5 O benzaldehyde diphenyl ketone (benzophenone) Chemical properties H Carbon atom of the carbonyl group is in the sp2-hybridization (flat fragment). Electrons of the double bond are strongly biased toward the more electronegative oxygen atom (C = O bond polar). Redistribution of charges in the carbonyl group has an effect on the polarity of the C-H bonds adjacent to the carbonyl group carbon atom (-position): H R C + C _ O H H a. Addition reactions involving C = O groups (nucleophilic addition SN) CH3 C O H X CH3 CH X a) Addition of hydrogen (reduction of aldehydes and ketones to primary and secondary alcohols) 131 O CH3 C t, Ni + H2 CH3 CH2 OH H t, Ni CH3 C CH3 + H2 CH3 CH CH3 OH O b) Addition of alcohols CH3 C O + HOR H CH3 CH OH OR (hemiacetal); OH CH3 C CH3 CH3 C CH3 + HOR OR O (acetal). c) Addition of sodium hydrosulfite (the reaction is used for separation of aldehydes and ketones from mixtures with other organic substances): O CH3 C + NaHSO3 H (HO-SO Na) 2 CH3 CH OH OSO2Na OH CH3 C CH3 CH3 C CH3 + NaHSO3 (HO-SO2Na) O OSO2Na. d) Addition of ammonia and ammonia derivatives CH3 CH3 CH3 C CH3 + H O 132 NHR CH3 C NH R OH _ CH3 C N R H2O àçî ì åòèí û (î ñí î âàí èÿ Ø èô ô à) azometines (Shiff bases) O H OH + HN NH2 H3C C H3C CH NH NH2 _ H H2O H3C CH N NH2 ãèäðàçî í hydrazone 2. Substitution reactions at -carbon atom O H3C C O + 3 Cl2 Cl3C C _ 3 HCl H H Iodoform test is a qualitative reaction for a carbonyl group: O H3C C I2, NaOH CHI3 + RCOONa èî äî ô î ðì (æåëòû é î ñàäî ê) R Iodoform (yellow precipitate) Since the primary and secondary alcohols can be oxidized by iodine to aldehydes and ketones, they also show the Iodoform test: H3C CH2 OH I2, NaOH CHI3 + HCOONa 3. Condensation reactions of aldehydes and ketones a) Aldol condensation: O H O + H2C C H3C C H OH H3C CH CH2 C H O H 133 CH3 H3C C CH2 C CH3 H3C C CH3 + H2C C CH3 O OH H O O b) Crotonic condensation: CH3 CH3 to H3C C CH2 C CH3 _ H3C C CH C CH3 H2O OH O O H3C C CH3 + H2C C CH3 O H O 4. Oxidation of aldehydes a) Silver mirror reaction (qualitative reaction for aldehydes): O CH3 C Ag2O (4NH3) t O CH3 C H или 2[Ag(NH ) ]OH 3 2 (аммиачный раствор оксида серебра) Ag + 3NH3 ONH4 (металлический налет на стенках пробирки) c) Reaction with copper(II) hydroxide (qualitative reaction for aldehydes): O t CH3 C 2Cu(OH)2 H (голубой осадок) O H C 4[Ag(NH3)2]OH H t O H C 4Cu(OH)2 H 134 O CH3 C Cu2O OH (красный осадок) (NH4)2CO3 Ag + 6NH3 t CO2 2Cu2O 5 Ketones can not be oxidized under the given conditions. CARBOXYLIC ACIDS Carboxylic acids compose a class of organic compounds containing a carboxyl functional group O . C OH In its structure, the carboxyl group can be associated with only one carbon atom (end group), so the isomerism of carboxylic acids is associated with the isomerism of the carbon skeleton or position of multiple bonds. Classification and nomenclature of carboxylic acids 1. Saturated monocarboxylic acids (general formula СnH2n+1COOH or CnH2nO2). Formula Н–СООН СН3–СООН СН3–СН2–СООН СН3–СН2–СН2–СООН CH3 CH COOH CH3 СН3–СН2–СН2–СН2–СООН С15Н31СООН С17Н35СООН Name (common system) methanoic (formic) ethanoic (acetic) propanoic butiric isobutiric valeric palmitic stearic 135 2. Unsaturated monocarboxylic acids (general formula СnH2n–1COOH or CnH2n–2O2) Formula Name (common system) СН2=СН–СООН acrylic CH2 C COOH metacrylic CH3 СН3–СН=СН–СООН С17Н33СООН crotonic oleic С17Н31СООН lynolic C O OH С17Н29СООН lynolenic C O OH С19Н31СООН araquidonic C O OH 3. Saturated dicarboxylic acids Formula Name (common system) НООС–СООН Ethanadioic (oxalic) НООС– СН2–СООН Propanedioic (malonic) НООС– (СН2)2–СООН Butanedioic (succinic) НООС– (СН2)3–СООН Pentanedioic (glutaric) НООС– (СН2)4–СООН Hexanedioic (adipic) 4. Unsaturated dicarboxylic acids 136 H H H C C HOOC COOH Maleic acid (cis-isomer) 5. Aromatic carboxylic acids Formula COOH C C HOOC H Fumaric acid (trans-isomer) Name COOH benzoic COOH COOH phthalic COOH isophthalic COOH COOH terephthalic COOH Polyfunctional acids 6. Hydroxyacids Formula * CH3 CH COOH Name lactic OH 137 malic * CH2 CH COOH HOOC OH OH HOOC citric C CH2 COOH CH2 COOH tartaric * * HOOC CH CH COOH OH OH salicylic COOH OH Oxyacids usually possess optic activity: COOH H Mirror COOH OH CH3 D- lactic acid OH HO H COOH D-tartaric H CH3 L- lactic acid COOH H HO COOH HO H COOH H H OH OH H OH COOH L-tartaric COOH mesotartaric Optically inactive because of the inner 138 symmetry plane Chemical properties The structure of the carboxyl group O H R C D H C C B A O H The reaction centers in the molecules of carboxylic acids are: A. OH of the carboxyl group. B. C-OH of the carboxyl group. C. R-COOH. D. C-H bond of the hydrogen atom in the -position A. Dissociatopn of the COOH group: RCOOH RCOO– + H+ (carboxylic acids are weak acids) The C-O bonds in the formed carboxylate anion are aligned, the negative charge is evenly distributed between the two oxygen atoms, therefore, the double C = O bond in carboxylic acids is inactive.: O _ R C O Acidic properties of carboxylic acids: 2СH3COOH + Fe (СH3COO)2Fe + H2; 2CH3COOН + MgO (CH3COO)2Mg + H2O; CH3COOH + KOH CH3COOK + H2O; 139 CH3COOH + NaHCO3 CH3COONa + CO2 + H2O. B. Elimination of the hydroxyl radical OH (nucleophilic substitution in carboxylic acids, acylation reaction) Esterification: O CH3 C + HO CH3 O H2SO4(конц.) CH3 C OH + H2O O CH3 (methyl acetate) Formation of acid halides: O O + PCl5 CH3 C CH3 C + POCl3 + HCl Cl OH Acetyl chloride O O + SOCl2 CH3 C CH3 C Formation of amides O t CH3 C + NH3 OH + SO2 + HCl Cl OH O CH3 C + H2O NH2 Formation of acid anhydrides: O O OH t , H2SO4(конц.) CH3 C O + H2O OH CH3 C CH3 C O O (acetic anhydride) CH3 C С. Decarboxylation 140 ление сплав CH4 + Na2CO3 t СН3СОСН3 + СаО + СО2 CH3COONa + NaOH (СН3СОО)2Са D. Radical reactions CH3–CH2–COOH + Сl2 CH3–CHCl–COOH + HCl Special properties of dicarboxylic acids НООС–СООН t СО2 + НСООН O O H2C C H2C C OH OH O H2C C t O + H2O H2C C O AMINO ACIDS AND PROTEINS Aminoacids are the molecular blocks of the molecular structure of the important and very complex class of compounds known as proteins. Natural aminoacids are -aminoacids where amino- and carboxylic groups occupy neighboring positions: 141 R CH COOH . NH2 Some examples of aminoacids are presented below: CH 2 COOH Aminoacetic acid (glycine, Glu) NH2 CH 3 CH COOH NH2 H3C CH CH2 NH2 O C OH HOOC CH 2 CH COOH NH2 NH 2 CO CH 2 CH COOH NH2 HOOC (CH 2 ) 2 CH COOH NH2 NH 2 CO (CH 2 ) 2 CH COOH NH2 -aminopropionic acid (alanine, Ala) 3-aminobutiric acid (β-butiric acid) Aspargic acid (Asp) Aspargine (Asn) Glutamic acid (Glt) Glutamine (Gln) HO CH 2 CH COOH NH2 Serine (Ser) HS CH 2 CH COOH NH2 Cysteine (Cys) CH 2 CH COOH NH2 142 Phenylalanine (Phe) HN Histidine (His) CH 2 CH COOH NH2 N Chemical properties Aminoacids are amphoteric compounds which combine the properties of weak organic acids and weak bases: H2N–CH2–COOH + NaOH H2N–CH2–COONa + H2O H2N–CH2–COOH + HCl Cl–H3N+–CH2–COOH nature: The solutions of amino acids are neutral because of its dipolar H2N–CH2–COOH H3N+–CH2–COO– (Zwitterion) Action at heat: -aminoacids O H N C O HO C CH R + H2N C OH NH2 R CH t CH R 2H2O + R CH C N H O O -aminoacids R CH CH2 COOH t NH3 + R CH CH COOH NH2 Other aminoacids H2N(CH2)nCOOH t (CH2)n C O NH 143 Proteins are products of condensation of amino acids of different nature: OH R1 R2 O H2N CH C NH CH C + H2N CH C H2N CH C O O O OH R1 R2 OH + H2O (dipeptide) O The C NH fragment is known as peptide (or amide) bond. The dipeptide can further react with amino acids, forming long polypeptide chains containing more than 40 amino acid residues, - proteins. They are also called polypeptides. The molar masses of proteins are in the range from 10,000 to several million au. The structure of natural proteins are 20 -amino acids. Some of them may be synthesized by the body (nonessential amino acids), and the other only comes with food (essential amino acids). The order of amino acids in protein molecules is called the primary structure of the protein. Due to intramolecular hydrogen bonds, protein molecules fold into spirals, forming a secondary structure of the protein. Helix, in turn, form the secondary coil (tertiary structure of the protein). Spirals can be combined in pairs by intermolecular hydrogen bonds, forming a quaternary structure. Depending on the shape of macromolecules distinguish globular (spherical) and fibrillar (filamentous) proteins. Under the influence of high temperatures, strong acids and alkalis, salts, toxic metals, radiation, quaternary, tertiary and secondary structure of proteins partially or completely destroyes in the process of protein denaturation. Qualitative reaction of peptides and proteins is the appearance of red-purple color when added the slurry of cupric hydroxide in an alkali to the protein solution (biuret reaction). 144 ESTERS, FATS AND OILS Esters are products of the reaction of alcohols with carboxylic acids (esterification) O H C + HO CH2 CH3 OH O реакция этерификации H C + H2O O CH2 CH3 ethylformiate The major chemical property of esters is reaction of hydrolysis (saponification) which is irreversible in alkaline media: O H C + H2O O CH2 CH3 H+ O H C O O H C + HO CH2 CH3 OH + NaOH O CH2 CH3 H C + HO CH2 CH3 ONa Esters of carboxylic acids are divided into three groups: 1. Fruity esters. These are esters of lower monocarboxylic acids and lower monovalent alcohols. They are liquids with a pleasant fruity and floral scent: Formula Name Scent 145 HCOOC2H5 ethylformiate rome HCOOCH2CH(CH3)2 isobutilformiate strawberry HCOOCH2C6H5 benzylformiate jasmin CH3COOC5H11 n-amylacetate pear CH3COOCH2CH2CH(CH3)2 isoamylacetate banana CH3COOC8H17 n-octylacetate orange C3H7COOC2H5 ethylbutirate pinapple Fruit esters are readily soluble in ethanol and diethyl ether, the solubility in water decreases with increasing molar mass. They are used as flavoring and fruit essences. 2. Waxes are esters of higher monobasic carboxylic acids and higher monohydric alcohols. They are a colorless solids, which are used for making candles, polishes and waxes floors, additives to soaps, lipsticks, etc. 3. Fats and oils are esters of higher (fatty) monobasic carboxylic acids and polyhydric alcohol - glycerol (triglycerides). Fats and oils are natural products. If the carboxylic acid is saturated (stearic, palmitic acids), the resulting ester is a fat (solid), and in the case of an unsaturated acid (oleic, linoleic, linolenic), this is the oil (liquid material). Typically, solid fats are of animal origin (beef and pork fat, butter), and liquid oils are vegetable. Fats are the feedstock for the manufacture of soap - sodium (solid soap) or potassium (liquid soap) salts of higher carboxylic acids: 146 CH2 O CO C17H35 CH O CO C17H35 + 3NaOH CH2 O CO C17H35 омыление CH OH + 3C17H35COONa CH2 OH глицерин тристеарат глицерина (твердый жир) glyceryl tristearate (fat) CH2 OH стеарат натрия (твердое мыло) glycerine sodium stearate (soap) Soaps have detergent properties as can emulsify fats and oils, i.e. convert them into fine droplets, which are wetted by water. Soaps can not be used in hard water since the calcium and magnesium ions form insoluble salts of fatty acids. Hydrogenolysis of fats is the process of adding hydrogen to unsaturated fatty acid residues that makes up the fat: CH2 O CO C17H33 CH O CO C17H33 + 3H2 CH2 O CO C17H33 òðèî ëåàò ãëèöåðèí à (ðàñòèòåëüí î å ì àñëî ) glyceryl trioleate (oil) t, Ni CH2 O CO C17H35 CH O CO C17H35 CH2 O CO C17H35 òðècòåàðàò ãëèöåðèí à (ì àðãàðèí ) glyceryl tristearate (solid fat) In this way, the conversion is carried out liquid fats (oils) in the solid, which is used in the production of food margarine - a substitute for butter. Fats are important as a food product, and characterized by the highest energy value. The technique fats are widely used for the manufacture of varnishes and oil paints. CARBOHYDRATES 147 Carbohydrates represent a class of organic compounds that contain groups of natural compounds of plant and animal origin. Most of the carbohydrate are of the general formula Cn(H2O)m. Carbohydrates include monosaccharides - polyhydroxy aldehydes (aldoses) and polyhydroxy ketones (ketoses), and their condensation products: disaccharides, and polysaccharides. According to the number of carbon atoms, monosaccharides are divided into tetroses (C4), pentose (C5), hexose (C6), and heptose (C7). In nature, the most common are pentoses and hexoses. Monosacharides Monosaccharides are the simpliest one-unit sugars. They contain asymmetric carbon atoms, so that possess a large number of stereoisomers, which are combined in pairs as optical antipodes - enantiomers. The number of stereoisomers N = 2n, where n is the number of asymmetric atoms in the molecule. For example, monosaccharides with four asymmetric carbon atoms exist in the form of 16 stereoisomers (enantiomers of 8 pairs of D- and L-series). Aldose D-series representing the natural products are listed below as Fischer projections: CHO CHO H OH HO H OH H CH2OH D-erythrose OH CH2OH D-threose aldotetroses 148 H CHO CHO H OH HO H OH H OH HO H OH H OH H CH2OH D-ribose CHO H CHO H OH HO H H HO H OH CH2OH CH2OH D-arabinose D-xylose aldopentoses CHO CHO OH CH2OH D-lyxose CHO OH HO H OH H OH HO H OH H OH H OH H OH H OH H OH H OH H OH CH2OH D-altrose CHO H OH HO H OH H HO H H OH CH2OH D-gulose HO H H H H H CH2OH D-allose CHO H CHO OH HO H H HO H CH2OH D-glucose CHO H CH2OH D-mannose CHO OH HO H OH HO H HO H H HO H HO H OH H OH CH2OH CH2OH D-idose D-galactose aldohexoses H OH CH2OH D-talose Diasteremerso are any combination of spatial isomers that are not a pair of optical antipodes. Pairs of diastereomers, differing in the configuration of one of more asymmetric atoms, are called epimers. For example, the epimers are 149 D-ribose and D-arabinose, as they only differ in the disposition of the substituents on the carbon atom in the 2nd position. Ketoses are monosaccharides containing a carbonyl (ketone) group. Some examples of ketoses are cited below: CH2OH CH2OH O HO O H H H OH HO CH2OH OH H CH2OH rybulose (product of photosynthesis in green plants) ketopentoses xylulose CH2OH CH2OH CH2OH O O O H H OH H H OH H OH HO H OH H OH H HO CH2OH fructose CH2OH psikose ketohexoses OH H OH CH2OH sorbose Formation of cyclic forms of monosaccharides The cyclization reaction is based on the ability of the alcohol and the aldehyde groups of the molecules to react with one another. The carbon chain containing sp3-carbons is bent, so a hydroxyl group at the fifth carbon atom comes close to the aldehyde group and is able to react with it. The formed during cyclization hydroxyl group at the 1st position can hold the 150 lower or upper position with respect to the plane of the ring, forming respectively - and -forms of the sacharide. For example, the reaction of cyclization of glucose can be presented as following: H 1 O C H C OH H H 2 OH HO H H OH H OH OH H O H O H 5 H CH2OH CH2OH CH2OH CH2OH 6 CH2OH H H H OH H HO 6 CH2OH OH OH 4 H H OH H OH OH H HO OH O C 3 H H HO 4C 5C O H OH OH C H 3 H H C2 O C 1 O H H OH H OH H H OH H OH OH -D-glucose -D-glucose (glucopyranose) (glucopyranose) The cyclization of monosaccharides is a reversible process, so the aldehyde group may be detected by a qualitative reaction (with Ag2O or Cu(OH)2). The hydroxyl group formed from the aldehyde group in the cyclization is called a glycoside (acetal) group. The term "anomers" refers to a pair of diastereomeric monosaccharide glycoside configuration differing atom in cyclic form, for example α-D- anomer and β-D-glucose. 151 The cyclization of fructose can lead to the formation of five-or sixmembered rings: CH2OH C O H O H H H OH OH OH H OH OH H C OH CH2OH OH CH2OH HO C H H C OH CH2OH H Fructopyranose H O HO CH2OH H fructofuranose Disaccharides Disaccharides are products of condensation of monosaccharides accompanied by an intermolecular dehydration. At formation of disaccharides, the monosaccharides may be same or different. One of the molecules participates in the reaction with its acetal OH group. The way of adjustment of another monosaccharide may be different. For example, in formation of sucrose (cane sugar), both OH fragments involving in the dehydration are of the acetal character: CH2OH O H H OH H OH OH H OH -D-glucose 152 H CH2OH H + H O HO CH2OH HO OH H fructose CH2OH O H H OH H CH2OH H H H O HO O OH H OH OH + H2O CH2OH H sucros e, С12Н22О11 (1-[ -D-fructofuranozyl]glucopyranozyde) -D- In this case, both the fragments of the disaccharide are in a stable cyclic form, and neither aldehyde nor ketone group may be formed. Such disaccharides are called non-reducing. Tregalose is another example of non-reducing disaccharides. It contains two glucose rests combined together through their acetal hydroxyl groups: Maltose is an example of a reducing disaccharide. Tho α-D glucose units are combined in the way that one of the glucoside fragment is not involved in condensation, and the cycle can open to produce an aldehyde group which possess reducing properties: 153 Another example of reducing disaccharides are lactose and cellobiose: (lactose) (cellobiose) Polysaccharides Polysaccharides are polycondensates of monosaccharides. Starch and cellulose are polysaccharides formed by chains of glucose fragments: nC6H12O6 (C6H10O5)n + nH2O. Starch is a mixture of polysaccharides derived from -D-glucose. It cans divided into two fractions: amylose (15 - 25%) and amylopectin (75 85%). In amylose, -D-glucose residues are linked with hydroxyl groups at the positions 1 and 4: O H H OH CH2OH CH2OH CH2OH H O H H OH H 154 OH O H H OH H H OH H H O O O H H H OH In amylopectin, the molecules of -D-glucose are linked through the hydroxyl groups at positions 1, 4 and 6: CH2OH O H H OH H H O H CH2OH O H H OH OH H CH2OH CH2 O H H OH H H H OH H H H O O O H O H H OH H OH OH Cellulose consists of residues of -D-glucose linked through the hydroxyl groups at positions 1 and 4: H OH O H H OH OH O H H OH H HO O H H H H H O H O H O H CH2OH CH2OH CH2OH H HO Reactions of polysaccharides: [C6H7O2(OH)3]n + 3nHO–NO2 cellulose nitric acid [C6H7O2(ONO2)3]n + 3nH2O nitrocellulose (explosive) [C6H7O2(OH)3]n + 3nHOОССН3 [C6H7O2(OОССН3)3]n + 3nH2O cellulose acetic acid acetylcellulose (synthetic fibers) 155 SULPHUR IN ORGANIC COMPOUNDS Organo-sulfur compounds are organic compounds that contain sulfur. Due to the fact that sulfur is analog of oxygen, it is able to replace it in the functional groups to form thiols (mercaptans), thioethers, thioaldehydes, thioacid. The name of the organic sulfur compounds usually appears prefix thio (thioalcohols thioaldehydes, monotio- and diotio carboxylic acids). Thioderivatives have stronger acidic properties and a very unpleasant odor; some are very toxic. BIOLOGICALLY IMPORTANT HETEROCYCLIC COMPOUNDS Heterocyclic compounds (heterocycles) are organic compounds containing rings whose composition comprises, together with the carbon atoms of other elements (nitrogen, oxygen, sulfur). By biological importance, nitrogen-containing heterocycles include five and six membered cyclic compound having one or two nitrogen atoms, and their condensed analogues are of a great significance. Pyrrole and its derivatives Pyrrole is a five- membered ring with one nitrogen atom. The lone pair of electrons of the heteroatom (nitrogen) is involved in conjugation with -electrons of the two double bonds, redistributing the electron 156 density. Thus, pyrrole readily undergoes electrophilic substitution on carbon atoms: Due to the fact that the pair of electrons of nitrogen is shifted towards the ring, the N-H bond is weakened and pyrrole exhibits acidic properties: + 2K 2 + H2 N K N H Pyrrole rings are part of the macrocyclic system – porphyrin. The porphyrin in the blood forms a complex with ferrous iron, which is called heme. Heme is in its turn connected with a coordination bond with globin (a protein consisting of the amino acids of high molecular weight). So constructed hemoglobin that is included in the composition of blood and performs the function of the oxygen transport in the body. CH CH2 H 3C HC N H H 3C N N CH CH H N HC HOOC CH2 CH N Fe N CH N HC CH3 CH CH2 CH2 HOOC CH2 porphyrin N HC CH2 CH3 hem Porfirin ring is also a constituent of chlorophyll - the green pigment in plants: 157 CH2 CH3 X N CH N Mg N CH N HC H2C CH C O H 3C HC CH COOCH3 H3C CH2 CH2 COOC20H39 chlorophyll (Х = СН3 or СНО) Condensation of the benzene ring to pyrrole molecule leads to the formation benzopyrrol or indole. Among the derivatives of indole there is an essential amino acid tryptophan, 3-(3-indolyl) -2-aminopropanoic acid: CH2 CH COOH NH2 N H N H indol triptophan Imidazole and its derivatives Imidazole is a five-membered ring with two nitrogen atoms in positions 1 and 3, one of the nitrogen atoms exhibits the basic properties,, and the other is acidic, so imidazole is amphoteric: N H+Cl- N + HCl N H 158 N H N N + NH3 + NaNH2 N H N Na The aminoacid histidine is one of the most important natural derivatives of imildazole. At elevated temperatures or under enzymatic decomposition histidine is decarboxylated and converted to histamine: N N H CH2 CH COOH t N N H NH2 histidine CH2 CH2 NH2 + CO2 histamine Pyridine and its derivatives Pyridine is the most important six-membered heterocycle having one nitrogen atom. Due to the fact that the pair of electrons of nitrogen does not participate in conjugation with the -electron ring, pyridine possesses basic properties, stronger than aniline, but weaker than aliphatic amines: + HCl N N+ H Cl Among pyridine derivatives there are nicotinic acid ( carboxypyridine), its amide (PP vitamin), pyridoxal, pyridoxol, pyridoxamine (B6 vitamin), nicotine, etc.: 159 CONH2 COOH N N Nicotinic acid nicotinamide CH2OH CHO HO H3C CH2OH N HO CH2OH H3C pyridoxal CH2NH2 HO CH2OH H3C N pyridoxol N pyridoxamin N N CH3 nicotine A condensed system of pyridine and benzene is quinoline. Pyridine ring in it more aromatic, than benzene, so the total -bond belongs to a greater extent to the pyridine ring. As a result, positions 5 and 8 of a quinoline increase the activity of electrophilic substitution reactions: 5 6 NO2 4 10 3 + 2HNO3 7 8 9 N H2SO4 + 2H2O 2 N 1 NO2 Quinoline moiety is a part of some drugs and quinoline alkaloids, of which the most important is quinine (antimalarian): 160 CH CH2 CH2 CH2 HO CH N H3CO N Pyrimidine and purine bases Pyrimidine is an example of the nitrogen-containing heterocyclic compound having two nitrogen atoms. Pyrimidine cycles are included in the nucleic acids associated with protein synthesis in cells. At the hydrolysis of nucleic acids, three important pyrimidine derivatives (pyrimidine bases) are formed. These are uracil (2,6-dihydroxypyrimidine), thymine (5-methyl-2,6dihydroxy-pyrimidine), and cytosine (6-amino-2-hydroxypyrimidine): 4 5 N 6 3 2 N 1 pyrimidine H3C N N NH2 OH OH OH N N N OH N OH uracyl thymine cytosine An important property of the pyrimidine bases is the keto-enol (lactim-lactam) tautomerism: 161 O OH NH N N N H OH lactyme (enol) O lactame (ketone) Along with the pyrimidine bases, the nucleic acids include purine pyrimidine condensed system and imidazole. Directly purine bases adenine (6-aminopurine) and guanine (2-amino-6-hydroxypurine). participate in the construction of nucleic acids 6 1 5 N 8 2 N 3 4 N9 H purine OH NH2 7 N N N N N H adenine N N H2N N N H guanine Purine and pyrimidine bases play an important role in the metabolism in the body. In conjunction with a monosaccharide (ribose or deoxyribose), and phosphoric acid, they form the nucleotides which make up the nucleic acids (DNA and RNA): Fragment of a polymeric chain of DNA 162 Purine derivatives are also included in the composition of alkaloids (caffeine, theophylline and theobromine), toxins (e.g., saxitoxin) and other related substances (uric acid): O H 3C O N N O CH3 N H3 C N O CH3 O N N N N H CH3 N HN O CH3 N N CH3 caffein theophylline theobromine (coffee and tea alkaloid) (tea alkaloid) (cacao alkaloid) List of citated literature 1. Neil D. Jespersen. Chemistry: The Molecular Nature of Matter: 7th Edition. Wiley, John & Sons, Incorporated, 2014. 163 2. Neil D. Jespersen. Student Solutions Manual to Accompany Chemistry: The Molecular Nature of Matter: 7th Edition. Wiley, John & Sons, Incorporated, 2014. 3. Rao C.N.R. University General Chemistry. Macmillan India Limited, 1997. 4. Bahl B.S., Bahl A. A textbook of Organic Chemistry. S Chand & Company Ltd., 1997. 164 Oльга Владимировна Ковальчукова, Насрин Намичемази, Русул Алабада Химия (конспект лекций). Для студентов 1 курса медицинского факультета специальности «Стоматология» (на английском языке) Зав. редакцией Т.О. Сергеева Техн. редактор И.М.Любавская Тематический план 2015 г., № _______________________________________________________ Подписано в печать Формат 60881/4. Ротапринтная печать. Усл. печ. л. 10,25. Усл. кр.-отт. 10,25. Уч.-изд. № . Цена договорная Издательство Российского университета дружбы народов 117923, ГСП-1, Москва, ул. Орджоникидзе,3 _________________________________________________________ Типография Российского университета дружбы народов 117923, ГСП-1, Москва, ул. Орджоникидзе,3 165