Chemistry: Composition of Matter & Atomic Models PDF

CHEMISTRY The composition of matter The subject of study in chemistry is matter. It is, in fact, the discipline that studies: - The structure and composition of matter - The transformations that matter undergoes - The energy involved in these transformations Matter is made up...

CHEMISTRY The composition of matter The subject of study in chemistry is matter. It is, in fact, the discipline that studies: - The structure and composition of matter - The transformations that matter undergoes - The energy involved in these transformations Matter is made up of atoms, tiny particles that can combine with one another in different ways. 1. Mixtures and substances Regarding chemical composition, matter can exist in the form of mixtures or substances. A mixture is a physical combination of two or more substances and has a variable composition. They are classified into: - Heterogeneous: formed from two or more phases; they do not have a uniform composition at all points of the sample, and the components are easily distinguishable. - Homogeneous: formed from a single phase; the composition is uniform at all points of the sample, and the components are indistinguishable. Homogeneous mixtures are also called solutions. A phase is a portion of a system in which the intensive physical properties are identical at every point or vary continuously, separated from the rest by limiting surfaces. By separating a mixture into its components using physical methods, pure substances are obtained. The choice of separation method depends on the physical state of the components and their chemical and physical properties. The most commonly used separation methods are: - Filtration: separates the components of a heterogeneous mixture consisting of a liquid and a solid. This is done by passing the mixture through an absorbent paper filter. - Distillation: separates the components of a mixture by exploiting their different boiling points. - Chromatography: separates the components of a mixture based on the different speeds at which they migrate through a supporting material under the influence of a solvent flow (eluate). - Centrifugation: exploits the difference in density of the components using centrifugal force. - Solvent extraction: exploits the solubility of one of the substances in the mixture in a solvent. A pure substance is a portion of matter that has a uniform, constant, and well-defined composition. It is made up of a single type of fundamental unit: an atom, a molecule, or the unit cell of a lattice. Based on the possibility of breaking down a substance into simpler components, two types of substances are distinguished: - Compounds: can be broken down because they are composed of different types of atoms bound together to form three-dimensional crystal lattices (NaCl) or molecules (H2O, NH3). - Elements (elementary substances): can not be broken down because they are made up of identical atoms. From this, we can understand that: an atom is the smallest fraction of matter that retains the chemical characteristics, but not the physical ones of the element; a molecule is the smallest portion of a pure substance that maintains its chemical and physical properties. 2. Early atomic models Evidence for the existence of atoms was provided (late 18th century) by the observation that matter involved in chemical transformations follows two laws: the law of conservation of mass (Lavoisier, 1783), which states that in a chemical reaction, the sum of the masses of the reactants is equal to the sum of the masses of the products, and the law of definite proportions (Proust, 1799), which states that in a pure substance, the elements that compose it are combined according to a defined and constant weight ratio. - Dalton's atomic theory (1802) Dalton's atomic theory is based on three postulates: 1: Elements are made of tiny particles called atoms. Atoms of an element are identical to each other and have the same chemical properties. Atoms of different elements are different and have different properties. 2: In chemical reactions, atoms retain their identity. 3: Atoms of different elements combine to form compounds. In a given compound, the relative number and type of atoms of each element are constant. The two previous laws can be explained in light of Dalton's theory. A consequence of his theory is the law of multiple proportions: when two atoms combine to form more than one compound, the weights of one that combine with a fixed weight of the other are in simple ratios expressible by small whole numbers ( 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑂 𝑖𝑛 𝐶𝑂2 32𝑔 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑂 𝑖𝑛 𝐶𝑂 = 16𝑔 = 2). - Rutherford model (1911) In the Rutherford model, the atom consists of a central nucleus made up of positively charged particles, called protons, and neutral particles, called neutrons, where almost all of the atom's mass resides. Surrounding the nucleus are negatively charged particles, called electrons, which move around the nucleus in a number equal to that of the protons. Therefore, the isolated atom is electrically neutral. Each atom is characterised by: - Z (atomic number): the number of protons contained in the nucleus - A (mass number): the number of neutrons and protons contained in the nucleus (e.g., 168O where Z=8 and A=16). All atoms with the same atomic number (Z) behave chemically in the same way and are classified as atoms of the same chemical element. Atomic radius = 10−10 m=1 Å Nuclear radius = 10−14 m=10−4 Å An atom can lose or gain one or more electrons (ionisation process), thus losing its electrical neutrality and transforming into an ion. A cation (positive ion) is an atom that has lost one or more outer electrons (e.g., H+); an anion (negative ion) is an atom that has gained one or more outer electrons (e.g., Cl−). However, the Rutherford model conflicts with classical electromagnetic theory, which posits that a moving electric charge, like an electron, should gradually lose energy, describing orbits of ever-decreasing radius until it falls into the nucleus. - Bohr model (1913) Bohr proposed a model in which the existence of stationary orbits is hypothesised, where electrons move without emitting energy. An electron belongs to a stationary orbit if the value of its angular momentum (mrv) is an integer ℎ −34 multiple of 2𝝅 (where ℎ = 6. 625 ∗ 10 𝐽 ∗ 𝑠 – Planck's constant), which means the following relationship holds: ℎ 𝑚𝑟𝑣 = 𝑛 ∗ 2𝝅 , where n takes integer values and is called the principal quantum number. According to his model: the energy of an electron increases with increasing n; the position of the electron depends on its energy content (+ farther = + energy); an atom can exchange energy with the outside only if one of its electrons transitions from one stationary orbit to another. If this occurs, the exchanged energy is equal to the difference in energy between the two states involved in the transition. Energy exchanges between an atom and its surrounding environment occur through the absorption (or emission) of a photon with 𝐸 = ℎ𝑣. - Wave model of the atom (1930) In this model, the atom consists of a nucleus containing protons and neutrons, and electrons. Here, however, the movement of electrons around the nucleus can only be represented using the concept of probability. The electron is confined to regions of space called orbitals, where it cannot be identified as a physical particle, behaving instead like a more or less dense electron cloud. Heisenberg's uncertainty principle states that it is impossible to simultaneously know the position and momentum of a particle. Consequently, for an electron, one cannot speak of a trajectory but only of a region of space where the probability of finding the electron at a given moment is greater than zero. An orbital is the region of space around the nucleus where there is a high probability of finding the given electron. Each orbital is identified by a mathematical function called the wave function ψ, which assigns probabilities of finding an electron with a certain energy in various regions of space around the nucleus. In this wave function, some numerical constants appear, known as quantum numbers. In general, it can be stated that: each orbital is determined by the three quantum numbers n, l, m (ms); each orbital can host a maximum of two electrons. 3. Electron arrangement in atoms The distribution of electrons in the various energy levels and sublevels of an atom is represented by the electronic configuration. The order of filling the orbitals is shown in the accompanying image. 4. Periodic table of elements Each chemical element is identified by a name and a symbol composed of one or two letters. The elements are arranged in increasing order of atomic number in horizontal rows, starting a new line when a new energy level begins to fill. Each horizontal row is called a period and corresponds to the filling of the orbitals of a level; each column is called a group. The elements in the same group have the same outer electron configuration, meaning they have the same number of electrons in the outermost energy level, known as the valence level, and the electrons residing there are called valence electrons. There are 7 periods, indicated by Arabic numerals, while the groups are indicated by Roman numerals. The table can also be divided into blocks: s, p, d, f. - Periodic properties of elements The chemical properties of atoms are determined by the electron configuration of the outermost shell. - Atomic radius: increases from top to bottom (groups) and increases from right to left (periods). - Ionisation energy: the energy required to remove an electron from an atom (cation). - Electron affinity: the energy released when a neutral atom gains an electron (anion). - Electronegativity: the tendency of an atom to attract bonding electrons. These three properties increase from left to right (periods) and increase from bottom to top (groups). - Classification of elements Based on chemical properties, elements can be classified into three groups: metals (to the left of the zigzag line – H is a non-metal); non-metals (to the right of the zigzag line); metalloids (straddling the line). Group IA: alkali metals; Group IIA: alkaline earth metals; Group VIIA: halogens; Group 0: noble gases. 5. Isotopes Isotopes are atoms that have the same atomic number (Z) but different mass numbers (A) because they contain a different number of neutrons. The isotopes of an element occupy the same place in the periodic table and have the same name. The only element whose isotopes have their own names is hydrogen: Protium (1 proton, 0 neutrons), Deuterium (1 proton, 1 neutron) and Tritium (1 proton, 2 neutrons). Since they have the same number of protons, and therefore the same number of electrons, the isotopes of a given element have the same chemical properties. 6. Atomic mass unit and mole Atomic mass unit (a.m.u.) = the amount of matter equal to 1/12 of the mass of an atom of the carbon-12 isotope −27 (126C), which is conventionally assigned a mass of 12. 1 𝑎. 𝑚. 𝑢. = 1. 67 ∗ 10 𝑘𝑔. Relative atomic weight (RA): the ratio of the absolute mass of the atom to the atomic mass unit. Molecular weight (MW): the sum of the atomic weights of the atoms that make up the molecule, each multiplied by its respective subscript. Example: 𝑃𝑀𝐻2𝑂 = 𝑃𝐴𝑂 + 2 ∗ 𝑃𝐴𝐻 = 16 + 2 ∗ 1 = 18 A mole is the quantity of matter that contains a number of elementary entities equal to the number of atoms 23 present in 12g of carbon-12. One mole of any substance contains 6. 02 ∗ 10 elementary units of that substance (Avogadro's number – NA​). From this definition, it follows that one mole of an element corresponds to the quantity of substance whose weight, expressed in grams, is numerically equal to its atomic weight. Example: one mole of 23 water (MW = 18) weighs 18g and contains 6, 02 ∗ 10 molecules of water. 𝑔 In light of the definition of a mole, it can be said that 𝑃𝑀𝐻2𝑂 = 18 𝑚𝑜𝑙. Using a formula, it is possible to calculate how many moles correspond to a given mass of substance (in grams) and vice versa. The number of moles (n) is calculated by dividing the given mass (m) by the mass of one mole of 𝑔 𝑚(𝑔) the substance considered, indicated as MW and expressed in 𝑚𝑜𝑙. Formula: 𝑛 = 𝑔 ​ 𝑃𝑀 ( 𝑚𝑜𝑙 ) - Avogadro's law Avogadro's law states that at equal conditions of pressure and temperature, equal volumes of different gases contain an equal number of molecules and therefore of moles. At 0°C and 1atm, one mole of any gas occupies 22.4L. 7. Chemical formulas The empirical formula indicates the type and number of atoms that make up a molecule, without showing how they are bonded together. It can be expressed as a minimum formula (CH₂O for glucose) and a molecular formula (C₆H₁₂O₆ for glucose). The structural formula indicates the spatial arrangement of atoms in the molecule, showing how they are bonded and the type of bond (CO₂: O = C = O). - Calculation of the empirical formula To calculate the empirical formula, you need to know the elements it consists of and the weight percentage of each element. For example, calculate the empirical formula of ethane knowing that its percentage composition is: C = 80% and H = 20%. This means that 100g of the compound contains 80g of C and 20g of H. You must find the number of moles by dividing the masses in grams by their respective molar masses: 18𝑔 ○ nC= 𝑔 = 6. 66 𝑚𝑜𝑙𝑒𝑠 12 𝑚𝑜𝑙 20𝑔 ○ nH= 𝑔 = 20 𝑚𝑜𝑙𝑒𝑠 1 𝑚𝑜𝑙 Next, divide the obtained values by the smallest value to get the simplest ratio: 6.66 ○ C= 6.66 =1 20 ○ H= 6.66 ≃ 3​ Thus, the empirical formula of ethane is: CH₃. - Calculation of the molecular formula To calculate the molecular formula, you need to know the empirical formula and the molecular weight, then multiply the subscripts from the empirical formula by a value equal to the ratio of the molecular weight (MW) to the formula weight of the empirical formula. For example, indicate how many H atoms are in the formula (NH4)2SO4​: 8 H atoms. - Calculation of the percent composition of a compound 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑋 𝑖𝑛 𝑜𝑛𝑒 𝑚𝑜𝑙𝑒 𝑜𝑓 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑∗100 To calculate percent composition, apply the following formula: %𝑋 = 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑡ℎ𝑒 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑 Example: calculate the percent composition of nitrous anhydride N2O3N_2O_3N2​O3​. ○ PAN = 14 ○ PAO = 16 ○ PMN2O3 = 2×14+3×16=28+48 = 76 28∗100 ○ %N = 76 ≃ 37% 48∗100 ○ %O = 76 ≃ 63% Chemical bonds Only noble gases exist in nature in a monatomic state. Other atoms are unstable and tend to bond with each other through chemical bonds to form elemental substances or compounds. 1. Valence electrons and Lewis symbols In the formation of chemical bonds, only the electrons, particularly valence electrons, are involved. Valence electrons of an element can be graphically represented by the Lewis symbol, which consists of the chemical symbol of the element surrounded by a number of dots equal to its number of valence electrons. 2. Octet rule The outer electron configuration of noble gases (s²p⁶) is called the octet: an extremely stable configuration. The octet rule states that atoms, when forming bonds, tend to achieve this electronic configuration by: sharing electrons (formation of a covalent bond) or losing or gaining electrons (formation of an ionic bond). This rule does not apply to transition elements. - Bond energy The formation of a chemical bond is a spontaneous phenomenon that occurs according to the physical principle whereby every system tends to reach a state of minimum potential energy, corresponding to maximum stability. A molecule forms if its energy is less than the total energy of the isolated atoms. When two atoms bond, a certain amount of energy is released, equal to the difference in energy between the initial and final states. The energy released during the formation of the bond, or required to break it, is called bond energy. The bond strength is proportional to the value of bond energy. 3. Covalent bond A covalent bond consists of a pair of electrons shared between two atoms; it forms between atoms whose difference in electronegativity is less than 1. 7. They share electrons to achieve an octet. For the bond to form, two atoms must come close enough together. The distance between the two nuclei at which an equilibrium is established between attractive forces (nucleus-electron) and repulsive forces (nucleus-nucleus and electron-electron) represents the bond length, measured in Ångströms (Å) and determined experimentally. The bond length is proportional to the size of the bonded atoms and inversely proportional to the bond energy.  - Simple and multiple covalent bonds Based on the number of shared pairs, the bond can be: single, double, or triple. The number of pairs of bonding electrons is called bond order. The greater the bond order, the stronger the force holding the atoms together, and therefore the higher the overall bond energy and the shorter the bond distance. - Molecular orbitals, σ bonding and 𝜋 bonding When a covalent bond forms, two atomic orbitals overlap and combine to form a molecular orbital, which has different energy and shape compared to the starting orbitals. A molecular orbital is the region occupied by a pair of bonding electrons surrounding the nuclei of the two bonded atoms. If the overlap is head-on, a σ covalent bond forms; if it is side-on, a 𝜋 covalent bond forms. The σ bond is stronger than the 𝜋 bond. A simple covalent bond is always of the σ type; a double bond consists of one σ bond and one 𝜋 bond; a triple bond consists of one σ bond and two 𝜋 bonds. - Polarity of the covalent bond Based on the difference in electronegativity, a covalent bond can be pure or polar. A pure covalent bond forms between atoms with equal or similar electronegativities (less than 0. 4) and is characterised by an equal sharing of the pair of bonding electrons. A polar covalent bond, on the other hand, forms between different atoms with a small difference in electronegativity (0. 4 − 1. 7) and is characterised by unequal sharing of the bonding electrons: the less electronegative atom acquires a partial positive charge (δ+), while the more electronegative atom acquires a partial negative charge (δ-). 4. Dative bond (coordination bond) This is a particular type of covalent bond, consisting of two electrons shared between two atoms. While in a covalent bond the two electrons of the shared pair are each provided by one atom, in a dative bond both electrons come from the same atom (the donor), which shares them with the other atom (the acceptor). The bond is indicated by an arrow pointing from the donor to the acceptor. 5. Hybridization and lone orbitals Hybridization involves the combination of the outer atomic orbitals of an atom, with different energies, resulting in an equal number of isoenergetic atomic orbitals called hybrid orbitals. Carbon has an outer electronic configuration of 2s² 2p². Having only two unpaired electrons, it should only be able to form two covalent bonds. However, in methane (CH₄), a carbon atom bonds with four hydrogen atoms through four identical covalent bonds. The formation of these four bonds is possible due to the creation of four isoenergetic hybrid orbitals, each containing one unpaired electron. - sp³ hybridization Occurs when one s orbital and three p orbitals of the same atom combine to form four isoenergetic hybrid sp³ orbitals directed toward the vertices of a regular tetrahedron, with angles of about 109°. - sp² hybridization In ammonia (NH₃), nitrogen is sp³ hybridised, and the bond angles are about 108°. It occurs when two p orbitals and one s orbital of the same atom combine to form three isoenergetic sp² hybrid orbitals arranged at 120° from each other. - sp hybridization In the BH₃ molecule, boron has three sp² hybrid orbitals, each containing one unpaired electron. It occurs when one p orbital and one s orbital of the same atom combine to form two isoenergetic sp hybrid orbitals arranged at 180°. In the BeH₂ molecule, Be is sp hybridised, and the bond angles are 180°. The spontaneity of the hybridization process arises from the fact that, in all cases, it leads to an increase in the stability of the molecules. Hybrid orbitals are atomic orbitals, not molecular orbitals. 6. Resonance and electron delocalization For some substances, it is possible to write more than one structural formula. This phenomenon is called resonance and occurs in substances where there is an extended 𝜋 electron system involving more than two atoms. The different possible formulas are called resonance structures and they are separated by double-headed arrows. The actual substance has a structure that is an average of the resonance forms and is called a resonance hybrid. The difference in energy between the resonance hybrid and the most stable resonance form is called resonance energy. 7. Molecular geometry – VSEPR theory The geometry of a molecule depends on the number and type of electron pairs surrounding the central atom. This geometry can be predicted based on the valence shell electron pair repulsion theory, known as VSEPR theory. This theory states that the valence electron pairs arrange themselves as far apart as possible to minimise electrostatic repulsion between them. To apply VSEPR theory, it is necessary to count the number of lone electron pairs and the number of atoms, thus determining the number of àshared pairs around the central atom. The total number of lone pairs plus shared pairs is called the steric number (SN). Different SN values correspond to different molecular geometries. 8. Polarity of molecules Molecules containing only pure covalent bonds are always nonpolar. Molecules formed by two atoms linked by a polar covalent bond are always polar. Molecules formed by more than two atoms linked by polar bonds are polar if the individual dipoles do not cancel each other out. For dipoles to cancel, they must have equal magnitude and opposite direction. 9. Ionic bond An ionic bond forms between two atoms with a high difference in electronegativity (greater than 1. 7) due to the transfer of one or more valence electrons from the less electronegative atom to the more electronegative one, resulting in the formation of two ions with opposite charges. In the case of ionic bonding, neither molecular orbitals nor molecules form. The force of the ionic bond is purely electrostatic. Rather than being composed of molecules, ionic compounds are made up of a collection of oppositely charged ions arranged to form a three-dimensional crystal lattice, where repulsive forces between like-charged ions are minimised, and attractive forces between oppositely charged ions are maximised. Ionic compounds have high melting and boiling points and conduct electricity in molten state and in aqueous solution. 10. Metallic Bond The bond between metal atoms in crystal lattices is neither ionic nor covalent. Metals tend to become cations because they easily lose their electrons. The lost electrons, shared among all ions and delocalized over an orbital extended throughout the metal, can move freely across the sample. The attraction between the cations of the crystal lattice and the delocalized valence electrons constitutes the metallic bond. Thus, the bond has both electrostatic and covalent character. 11. Intermolecular Bonds Molecules in a solid or liquid interact with one another through intermolecular bonds: attractive forces that are much weaker than the bonds that hold atoms together within molecules. These include dipole-dipole interactions, London forces (Van der Waals forces), hydrogen bonding, and ion-dipole interactions. When an ionic compound melts, the crystal lattice breaks. When a molecular compound melts, dispersion forces, dipole-dipole interactions, or hydrogen bonds are broken. - Dipole-dipole interactions These are attractive forces that arise between polar molecules. Such molecules behave like spontaneous and permanent electric dipoles, attracting each other by orienting with the positive end of one dipole close to the negative end of another dipole. - Dispersion forces These are extremely weak attractive forces related to the formation of temporary dipoles caused by the movement of electrons around the nucleus. Temporary dipoles interact with the electron clouds of nearby molecules, polarising them, creating induced dipoles, and establishing weak attractive forces. Dispersion forces occur between all types of molecules. The intensity of dispersion forces increases with molecular weight, surface area, and the number of electrons present in the molecule. 12. Hydrogen bond This is an electrostatic interaction between a hydrogen atom covalently bonded to a very electronegative atom (F, O, N) and the lone pair of a very electronegative atom (F, O, N) in an adjacent molecule (intermolecular) or in the same molecule (intramolecular). It is the strongest of the intermolecular attractive forces. Molecules connected by hydrogen bonds have boiling points higher than those of molecules of similar molecular mass that do not form hydrogen bonds. The presence of hydrogen bonding in water determines its characteristics: a decrease in density when transitioning from liquid to solid, and an increase in volume during the transition from liquid to solid. 13. Allotropism and polymorphism Allotropism: when an element exhibits two forms that differ in molecular structure or in how the atoms are bonded. These two forms have different physical and chemical characteristics (O₂ and O₃). Polymorphism: when a substance exists in forms that differ only in crystalline structure (sulphur). Inorganic Compounds and Nomenclature 1. Oxidation number Calculating the oxidation number of an atom is important to understand the processes occurring during chemical reactions. The oxidation number (O.N.) corresponds to the charge that the atom would assume if the electron pairs in bonds were assigned to the more electronegative atom involved. Rules for calculating the oxidation number: - All substances in their elemental state have an oxidation number of zero. - The oxidation number of a monatomic ion is equal to the charge of the ion. - The sum of the oxidation numbers of all atoms in a neutral molecule is equal to 0, while in a polyatomic ion, it is equal to the charge of the ion. - The O.N. of hydrogen is +1, except in compounds with metals, where it is -1. - In most compounds, the O.N. of oxygen is -2, but not in all cases. - The O.N. of group IA elements is +1; group IIA elements is +2; group IIIA elements is +3. 2. Binary compounds These are formed from only two different elements (CaCl₂, H₂O). To write the formula for a binary compound, a specific rule is followed: the cation (the less electronegative element) must be written before the anion (the more electronegative element). Regarding nomenclature, a binary compound is named after the more electronegative element, which is given the suffix -ide, followed by the name of the less electronegative element (e.g., calcium chloride). A metal may often exhibit two different oxidation numbers. In this case, it is necessary to distinguish the different compounds that a given metal can form with the same non-metal. Element Salt Traditional nomenclature Stock notation IUPAC nomenclature Copper Cu₂S Cuprous sulphide Copper (I) sulphide Dicopper monosulfide Cu₂S Cupric sulphide Copper (II) Copper monosulfide sulphide Iron FeCl₂ Ferrous chloride Iron (II) chloride Iron dichloride FeCl3 Ferric chloride Iron (III) chloride Iron trichloride - Basic and acidic oxides Oxides are binary compounds formed by the combination of one of several elements with oxygen. The formula is written by placing the symbol of the element with which it is combined before the symbol for oxygen and assigning appropriate subscripts based on the oxidation number of the element (O.N. of oxygen = -2). Basic oxides are binary ionic compounds formed from a metallic cation (Mx+) and the oxide ion (O²⁻). Formula Traditional nomenclature Stock notation IUPAC nomenclature FeO Ferrous oxide Iron (II) oxide Iron monoxide Fe₂O₃ Ferric oxide Iron (III) oxide Diiron trioxide Acidic oxides are binary compounds formed from a non-metal and oxygen. Formula Oxidation number Traditional nomenclature IUPAC nomenclature N₂O₃ +3 Nitrous anhydride Dinitrogen trioxide N₂O₅ +5 Nitric anhydride Dinitrogen pentoxide - Hydracids These are binary compounds formed from hydrogen (H) and one of the following non-metals: S, F, Cl, Br, I. In their pure state, these substances are molecular compounds; however, when dissolved in water, they behave as acids, releasing H⁺ ions. ○ HCl (gas): Hydrogen chloride ○ HCl (aqueous): Hydrochloric acid 3. Ternary compounds These are compounds formed by the combination of three elements. - Hydroxides (bases) These are ternary ionic compounds formed from a metallic cation (Mx+) and as many hydroxide ions (OH⁻) as needed to neutralise the charge of the cation. General Formula: M(OH)x. They are prepared by reacting basic oxides with water: CaO + H₂O → Ca(OH)₂. When dissolved in water, those that are water-soluble dissociate and behave as bases, releasing OH⁻ ions. Formula Traditional nomenclature IUPAC nomenclature Fe(OH)₂ Ferrous hydroxide Iron dihydroxide Fe(OH)₃ Ferric hydroxide Iron trihydroxide - Oxyacids (oxygen acids) These are ternary molecular compounds formed from hydrogen, a non-metal or a transition metal with another oxidation state, and oxygen, written in that order. General Formula: HxNMOy. The name derives from that of the corresponding anhydrides, replacing the word "anhydride" with "acid." When dissolved in water, they ionise and dissociate, releasing H⁺ ions. Formula Nomenclature H₂CO₃ Carbonic acid HNO₂ Nitrous acid HNO₃ Nitric acid H₂SO₃ Sulphurous acid H₂SO₄ Sulphuric acid Monoprotic acids are those that can donate a single H⁺ ion, forming one type of anion. Polyprotic acids are those that can donate more than one H⁺ ion, forming different types of anions. ○ HCl (Hydrochloric acid) → Cl⁻ (Chloride ion) ○ HClO₂ (Chlorous acid) → ClO₂⁻ (Chlorite ion) ○ HClO₃ (Chloric acid) → ClO₃⁻ (Chlorate ion) ○ CO₃²⁻ (Carbonate ion) ○ HCO₃⁻ (Bicarbonate ion) - Salts These are ionic compounds derived from acids by total (neutral salts) or partial (acid salts) replacement of hydrogen atoms with one or more metallic cations. Neutral salts are binary compounds, while acid salts are ternary compounds if derived from hydracids, and quaternary if derived from oxyacids. To write the formula, the cation is indicated first, followed by the anion. The name of the compound is obtained by stating the name of the anion first and then that of the cation. Example: LiCl → Lithium chloride. To derive the formula of a salt, one must identify the ions it is composed of and then construct the formula by balancing the charges of the ions. Example: Aluminium sulphate. Cation: Al³⁺. The suffix -ate indicates that sulphur is present with the highest oxidation number (+6). Sulphate ion: SO₄²⁻. The least common multiple is 6, so dividing this number by the valence of each ion yields the formula: Al₂(SO₄)₃. Salts can be prepared in many ways: ○ acid + base → salt + water (H₂SO₄ + Ca(OH)₂ → CaSO₄ + 2H₂O): neutralisation reaction. ○ acid + metal → salt + hydrogen (H₂SO₄ + Zn → ZnSO₄ + H₂). ○ anhydride + oxide → salt (CO₂ + CaO → CaCO₃).  ○ anhydride + hydroxide → salt + water (CO₂ + 2NaOH → Na₂CO₃ + H₂O). ○ acid + Oxide → salt + water (2HCl + Na₂O → 2NaCl + H₂O). ○ acid₁ + Salt₁ → salt₂ + acid₂ (2HCl + CaCO₃ → CaCl₂ + H₂CO₃). ○ salt₁ + Salt₂ → salt₃ + salt₄ (BaCl₂ + Na₂SO₄ → BaSO₄ + 2NaCl). - Preparation reactions of main inorganic compounds ○ metal + oxygen → oxide ○ oxide + water → hydroxide ○ non-metal + oxygen → anhydride ○ non-metal + hydrogen → hydracid ○ anhydride + water → oxyacid ○ hydroxide + hydracid → salt + water 4. Common elements: properties and main compounds - Hydrogen The simplest and most abundant chemical element; characterised by intermediate electronegativity, it forms compounds with many elements; can have oxidation states of +1, -1, or 0; it is a gas consisting of diatomic molecules (H₂). - Group IA (1): alkali metals Outer electron configuration: ns¹; low electronegativity, so they form monovalent positive ions and produce ionic compounds; oxidation state: +1; in elemental form, they are ductile and malleable solids with low melting and boiling points; the oxides of alkali metals are basic and react with water to form hydroxides. Lithium – Li (3), Sodium – Na (11), Potassium – K (19). - Group IIA (2): alkaline earth metals Outer electron configuration: ns²; low electronegativity, so they form divalent positive ions and produce ionic compounds; oxidation state: +2; in elemental form, they are solids; their oxides are basic and react with water to form hydroxides. Magnesium – Mg (12), Calcium – Ca (20), Barium – Ba (56). - Transition metals The outer electron configuration is characterised by filling d orbitals; they are hard solids (except mercury) with high melting points; low electronegativity, so they tend to transform into cations, forming ionic compounds. Iron – Fe (26), Copper – Cu (29), Zinc – Zn (30). - Group IIIA (13) Outer electron configuration: ns² np¹; oxidation state: +3; consists of elements with very different chemical characteristics. Boron – B (5), Aluminum – Al (13). - Group IVA (14) Outer electron configuration: ns² np²; they have very different electronegativity values and chemical characteristics; in elemental form, they are all solids. Carbon – C (6), Silicon – Si (14), Lead – Pb (82). - Group VIA (16) Outer electron configuration: ns² np³; Nitrogen – N (7), Phosphorus – P (15), Arsenic – As (33). Outer electron configuration: ns² np⁴; they are non-metals characterised by high electronegativity values; the most common oxidation state is -2. Oxygen – O (8), Sulphur – S (16). - Group VIIA (17): halogens Outer electron configuration: ns² np⁵; non-metals characterised by very high electronegativity; the most common oxidation state is -1; they tend to gain electrons, forming monovalent anions; their oxides react with water to form corresponding oxyacids; in elemental form, they consist of diatomic molecules; they form compounds with hydrogen known as halides. Fluorine – F (9), Chlorine – Cl (17), Bromine – Br (35), Iodine – I (53). - Group VIIIA (18): noble gases Outer electron configuration: ns² np⁶, except for helium (1s¹); they are stable; very high ionisation energy; do not form compounds; in elemental form, they exist as monoatomic gases. Helium – He (2), Neon – Ne (10), Argon – Ar (18), Krypton – Kr (36), Xenon – Xe (54). Chemical reactions 1. Definitions A chemical reaction is a transformation in which chemical bonds are broken and formed, and a certain amount of energy is exchanged. To write an equation in the most complete form, it is necessary to indicate the physical state of the reactants and products: (g): gas, (l): liquid, (s): solid, (aq): in aqueous solution. 2. Balancing chemical reactions In chemical reactions, the mass of the reactants must equal the mass of the products; therefore, an equation must be balanced. To balance an equation, coefficients are placed before the reactants and products to respect the law of conservation of mass.  3. Classification of chemical reactions There are several types of chemical reactions: - Synthesis reaction: two or more substances react to form a single substance (H₂ + I₂ → 2HI). - Decomposition reaction: a single compound breaks down into two or more substances (CaCO₃ → CaO + CO₂). - Dissociation reaction: a compound dissociates, releasing ions (MgCl₂ → Mg²⁺ + 2Cl⁻). - Ionisation reaction: a molecular compound reacts with water to form positive and negative ions (HCl + H₂O → H₃O⁺ + Cl⁻). - Neutralisation reaction: acid + base → salt + water (NaOH + HNO₃ → NaNO₃ + H₂O). - Substitution reaction: atoms of one element replace atoms of another in a compound (Fe + 2HCl → FeCl₂ + H₂). - Double exchange reaction: exchange of ions between two compounds (BaCl₂ + Na₂SO₄ → BaSO₄ + 2NaCl). - Condensation reaction: two molecules combine with the elimination of a water molecule (CH₃OH + CH₃OH → CH₃OCH₃ + H₂O). - Hydrolysis reaction: a molecule is broken down by the addition of water, forming two molecules (CH₃OCH₃ + H₂O → CH₃OH + CH₃OH). - Addition reaction: typical of organic compounds; a new element is added to the starting molecule (CH₃-CH=CH-CH₃ + H₂ → CH₃-CH₂-CH₂-CH₃). - Oxidation-reduction reaction: transfer of electrons (2AgNO₃ + Cu → Cu(NO₃)₂ + 2Ag). - Combustion reaction: the fuel oxidises in the presence of an oxidizer, usually oxygen, releasing thermal energy (CH₄ + 2O₂ → CO₂ + 2H₂O + heat). Often, a reaction can belong to multiple classes among those listed. 4. Ionic Equations Molecular equation: CaCl₂ + 2AgNO₃ → Ca(NO₃)₂ + 2AgCl Ionic equation: Ca²⁺ + 2Cl⁻ + 2Ag⁺ + 2NO₃⁻ → Ca²⁺ + 2NO₃⁻ + 2AgCl In this reaction, the ions Ca²⁺ and NO₃⁻ do not participate in the reaction and are therefore called spectator ions. The ionic equation can also be written as a net ionic equation, omitting the spectator ions and showing only those that actually participate in the reaction: Cl⁻ + Ag⁺ → AgCl. 5. Redox reactions (oxidation-reduction reactions) A redox reaction can be considered the sum of two half-reactions: oxidation and reduction. Oxidation: a chemical species is oxidised when it loses electrons and thus increases its oxidation state (N.O.). Reduction: a chemical species is reduced when it gains electrons and thus decreases its oxidation state. The species that is oxidised donates electrons to the species that is reduced, making it the reducing agent; the species that is reduced gains electrons from the species that is oxidised, making it the oxidising agent. - The electric battery A device that transforms the chemical energy released by an exothermic oxidation-reduction reaction into electrical energy. The anode is the negative electrode where the oxidation half-reaction occurs; the cathode is the positive electrode where the reduction half-reaction occurs. - Electrolysis An electrolytic cell is a device in which electrical energy is transformed into chemical energy. The supplied electrical energy allows an endothermic oxidation-reduction reaction to occur. Similar to a battery: oxidation occurs at the anode and reduction at the cathode, but in reverse, the anode is the positive electrode and the cathode is the negative one. 6. Stoichiometry Stoichiometry studies the mass relationships between elements in compounds and the mass relationships between reactants and products in a chemical reaction. All stoichiometric calculations are based on the observation of the balanced chemical equation, which can be read in four different ways: number of elementary entities; number of moles; mass; volume; only if reactants and products are gases and the reaction occurs under standard conditions. - Stoichiometric calculations Mole-mole: calculate the number of moles of NH₃ produced from the reaction of 6.33 moles of H₂. Reaction: N₂ + 3H₂ → 2NH₃. 3 𝑚𝑜𝑙 𝐻2 : 2 𝑚𝑜𝑙 𝑁𝐻3 = 6. 33 𝑚𝑜𝑙 𝐻2: 𝑋 2∗6.33 𝑋 = 3 = 4. 22 𝑚𝑜𝑙 𝑁𝐻3 Mole-mass: calculate the grams of NH₃ (PM = 17 g/mol) obtained from 2.3 moles of N₂. Reaction: N₂ + 3H₂ → 2NH₃. 1 𝑚𝑜𝑙 𝑁2: 2 𝑚𝑜𝑙 𝑁𝐻3 = 2. 3 𝑚𝑜𝑙 𝑁2: 𝑋 2∗2.3 𝑔 𝑋 = 1 = 4. 6 𝑚𝑜𝑙 𝑁𝐻3 → 𝑚 = 𝑛 ∗ 𝑃𝑀 = 4. 6 𝑚𝑜𝑙 ∗ 1. 7 𝑚𝑜𝑙 = 78. 2 𝑔 𝑁𝐻3 Solutions 1. Definitions A solution is a homogeneous mixture composed of at least two components: solvent and solute, whose quantities can be varied continuously. The solvent is the component present in the greater amount and exists in the same state of aggregation as the solution. The solute is the substance dissolved by the solvent and is present in a smaller quantity within the solution. 2. The dissolution process The process of dispersing solute particles among the solvent molecules is called dissolution or solvation. When the solvent is water, the process is referred to as hydration. - Solubility The amount of solute that can dissolve in a solvent is not unlimited. A saturated solution is one that contains the maximum amount of a given solute that a given volume of solvent can dissolve. The solubility of a solute is the concentration of the saturated solution. It is usually expressed in moles per litre but can also be expressed in grams per litre or in other forms of concentration. A solution with a relatively low concentration of solute is called dilute, while one with a relatively high concentration is called concentrated. These terms do not have a precise quantitative meaning. - Factors affecting solubility Substances in which particles are held together by similar intermolecular forces tend to be soluble in one another. Generally, it can be stated that "like dissolves like". The solubility of substances for which dissolution is an endothermic process increases with increasing temperature, according to Le Chatelier's principle. Conversely, the solubility of substances for which dissolution is an exothermic process decreases with increasing temperature. Pressure influences the solubility of gases in liquids. The solubility of a gas in a liquid is proportional to the partial pressure of the gas above the solution. - Factors affecting the rate of dissolution Agitating the solution increases the rate of dissolution of the solute in the solvent. Increasing the temperature raises the rate of the dissolution reaction. The greater the subdivision of the solute, the larger the contact surface between the solvent and the solute, thus increasing the rate of the dissolution reaction. - Aqueous solutions and solubility rules Water is an excellent solvent: it is a dipole and can form hydrogen bonds or ion-dipole bonds with the solute. Non-polar covalent solids and metallic solids are insoluble in water. The solubility of a molecular solid in water depends on its polarity and the ability to form bonds. Ionic solids are generally soluble in water, but not all of them. 3. Concentration of solutions The concentration of a solution expresses the ratio between the amount of solute and the amount of solution or solvent. Concentration can be expressed in different units of measurement: Percent composition by weight/weight (% w/w): indicates the number of grams of solute in 100 g of solution. 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑔) %𝑤/𝑤 = 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝑔) ∗ 100 Percent composition by weight/volume (% p/v): indicates the grams of solute in 100 mL of solution. 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑔) %𝑤/𝑣 = 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝑚𝐿) ∗ 100 𝑛 (𝑚𝑜𝑙) Molarity (M): the number of moles of solute contained in one litre of solution. 𝑀 = 𝑉 (𝐿) 𝑛 (𝑚𝑜𝑙) Molality (m): the number of moles of solute dissolved in 1 kg of solvent. 𝑚 = 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 (𝑘𝑔) Mole fraction (χ): the ratio of the number of moles of a component to the total number of moles of all 𝑛 𝑐𝑜𝑚𝑝𝑜𝑛𝑒𝑛𝑡 𝐴 (𝑚𝑜𝑙) components in the solution. χ𝐴 = 𝑛 𝑡𝑜𝑡𝑎𝑙 (𝑚𝑜𝑙) 𝑛𝑒𝑞 (𝑒𝑞) Normality (N): the number of equivalents of solute contained in 1 L of solution. 𝑁 = 𝑉 (𝐿) - Chemical equivalent and equivalent mass The chemical equivalent is a unit of quantity of matter whose definition depends on the type of substance considered and the reaction in which it is involved. One equivalent of a chemical species always reacts with one equivalent of another, yielding one equivalent product. The equivalent is a variable quantity; the number of equivalents is always an integer multiple of the number of moles.  𝑔 The mass in grams of one equivalent is called the equivalent weight (EW) or equivalent mass (EM) [ 𝑒𝑞 ]. 𝑔 𝑔 𝑃𝑀 ( 𝑚𝑜𝑙 ) 𝑚 (𝑔) Formulas: 𝐸𝑊 ( 𝑒𝑞 ) = 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑒𝑞𝑢𝑖𝑣𝑎𝑙𝑒𝑛𝑡𝑠 𝑐𝑜𝑛𝑡𝑎𝑖𝑛𝑒𝑑 𝑖𝑛 𝑜𝑛𝑒 𝑚𝑜𝑙𝑒 𝑜𝑓 𝑠𝑢𝑏𝑠𝑡𝑎𝑛𝑐𝑒 ; 𝑛𝑒𝑞 = 𝑔 𝐸𝑊 ( 𝑒𝑞 ) When diluting a solution, the number of moles of solute does not change; therefore, it holds that: 𝑀𝑐𝑜𝑛𝑐 ∗ 𝑉𝑐𝑜𝑛𝑐 = 𝑀𝑑𝑖𝑙 ∗ 𝑉𝑑𝑖𝑙. Where M is the molarity and V is the volume. 4. Ionisation and ionic dissociation Solutes that release ions when dissolved in water make the resulting solution conductive and are called electrolytes, while their solutions are referred to as electrolytic solutions. Ionic dissociation (electrolytic) refers to the process of separating the positive and negative ions present in an ionic compound, caused by water. Ionisation refers to the process of forming positive and negative ions operated by water on a polar molecular compound. Thus, water transforms the molecular compound into an ionic compound, which then dissociates. A strong electrolyte is a substance that completely dissociates in water; a weak electrolyte is a substance that partially dissociates in water. Soluble salts, strong acids, and strong bases are strong electrolytes; weak acids and weak bases are weak electrolytes. 5. Colligative properties of solutions When a non-volatile solute is added to a solvent, the physical properties of the solution differ from those of pure solvent. Colligative properties are properties of solutions of a non-volatile solute in a volatile solvent that depend solely on concentration, i.e., the total number of solute particles in solution (molecules or ions) and not on their nature. These include: lowering of vapour pressure; boiling point elevation; freezing point depression; and osmotic pressure phenomenon. - Lowering of vapour pressure: Raoult's law According to Raoult's law, the vapour pressure of an ideal solution is equal to the sum of the vapour pressures of the components, each multiplied by its respective mole fraction in the liquid phase. Consequently, the vapour pressure of a solution containing a non-volatile solute is always lower than that of the pure solvent. - Boiling point elevation and freezing point depression The addition of a non-volatile solute to a solvent results in a solution whose boiling point is higher (boiling point elevation) and whose freezing point is lower (freezing point depression) than that of the pure solvent. The magnitude of the effect is proportional to the molal concentration of the solution. - Osmosis and osmotic pressure Separating two solutions of different concentrations with a semipermeable membrane, which is permeable only to the solvent and not to the solute, results in the phenomenon of osmosis. This consists of the net movement of solvent through the membrane from the more dilute solution to the more concentrated one, until an equilibrium situation is reached. The result is the rise in the level of the more concentrated solution compared to the more dilute one. The pressure that must be applied to the more concentrated solution to bring it back to the level of the more dilute one is called osmotic pressure (π). The osmotic pressure of a solution in which the solute is a 𝑛 nonelectrolyte is calculated using the relationship: π ∗ 𝑉 = 𝑛 ∗ 𝑅𝑇 → π = 𝑉 ∗ 𝑅𝑇 = 𝑀 ∗ 𝑅𝑇. Comparing the osmotic pressures of two solutions: if they are equal, the two solutions are isotonic with respect to each other; if the osmotic pressure of A is greater than that of B, solution A is hypertonic compared to solution B, while solution B is hypotonic compared to solution A. - Colligative properties in electrolytic solutions To calculate the effect on colligative properties induced by the presence of an electrolyte, it is necessary to consider the total number of particles into which the electrolyte dissociates, introducing Van’t Hoff's factor (𝑖) into the formula. To calculate osmotic pressure, the formula π = 𝑀 ∗ 𝑅𝑇 ∗ 𝑖 can be used, or osmolarity can be utilised: an expression of the concentration of a solution that emphasises the number of solute particles present in the solution. It expresses the number of osmoles of solute contained in one litre of solution, where the number of osmoles is calculated by multiplying the number of moles by the coefficient i: 𝑂𝑆𝑀 = 𝑀 ∗ 𝐼, 𝑛𝑜𝑠𝑚 = 𝑛 ∗ 𝑖. 6. Suspensions and colloids Heterogeneous mixtures consist of multiple phases dispersed in one another, but distinguishable at least under a microscope. Depending on the phases in which the components are found, they can be distinguished as: - Aerosol: an heterogeneous mixture consisting of a solid or liquid phase dispersed in a gas phase (fog); - Emulsion: a mixture of two liquid phases dispersed in one another (oil and vinegar); - Suspension: a system consisting of a solid phase dispersed in a liquid phase (salt in oil); - Colloids: mixtures in which the particles have sizes intermediate between those found in solutions and those of coarse suspensions (pigments used for paints). Thermodynamics, kinetics of reactions and chemical equilibrium  1. Heat of reaction and enthalpy change Chemical reactions that release heat are called exothermic, while those that absorb heat are endothermic. To express the thermal changes accompanying a reaction, a state function called enthalpy (𝐻) is used, which indicates the thermal content of the system and can be imagined as energy stored in the form of chemical bonds. The heat (𝑄) absorbed or released by the system at constant pressure, known as the heat of reaction, is equal to the change in enthalpy (∆𝐻) of the reaction. Formula: ∆𝐻 = 𝐻𝑝𝑟𝑜𝑑 − 𝐻𝑟𝑒𝑎𝑐 = 𝑄. In an exothermic reaction, 𝐻𝑝𝑟𝑜𝑑 < 𝐻𝑟𝑒𝑎𝑐, so ∆𝐻 < 0; in an endothermic reaction, 𝐻𝑝𝑟𝑜𝑑 > 𝐻𝑟𝑒𝑎𝑐, so ∆𝐻 > 0. Heat acts as a reactant in endothermic reactions and as a product in exothermic ones, and thus should be indicated in the reaction equation. An equation that includes the heat data is called a thermochemical reaction. Hess's Law: since enthalpy is a state function, ΔH of a reaction depends only on the initial and final states and thus has a value independent of whether it occurs in a single stage or through a series of intermediate steps. 2. Spontaneity of chemical reactions and free energy Often, exothermic reactions are spontaneous. The spontaneity of a reaction depends on: the change in enthalpy (∆𝐻), temperature (𝑇), and change in entropy of the system (∆𝑆). There is a relationship among the three parameters expressed by Gibbs's equation, which indicates the behaviour of free energy (∆𝐺): ∆𝐺 = ∆𝐻 − 𝑇∆𝑆 If ∆𝐺 < 0, the reaction is spontaneous. A spontaneous reaction is called exergonic. If ∆𝐺 > 0, the reaction is nonspontaneous; in this case, it is called endergonic. If ∆𝐺 = 0, the reaction is at equilibrium. ∆𝐻 ∆𝑆 ∆𝐺 - + - Spontaneous at all temperatures + - + Non-spontaneous at all temperatures + + +/- Spontaneous only at high temperatures - - +/- Spontaneous only at low temperatures 3. Chemical kinetics This is the branch of chemistry that deals with the rates of chemical reactions and the factors that influence them. It is based on the collision theory: for the reaction A + B → C + D, for A and B to react, it is necessary that: - A and B collide: the collision theory states that the rate of a reaction is proportional to the number of collisions per second between the reacting molecules; - The collision is effective energetically and spatially: for the collision to be effective and the bonds between the reacting particles to break, they must not only possess sufficient kinetic energy but also collide with the correct mutual orientation. There is a threshold energy below which the collision between A and B does not lead to the formation of products: the activation energy (Ea), which corresponds to the minimum energy required for the formation of the activated complex. This is an aggregate of atoms in an infinitesimally short lifetime during which old bonds are breaking and new bonds are forming. The moment in the reaction when the activated complex forms is called the transition state. - Reaction rate Denoting the molar concentration of the reactant species as [A] and the reaction rate as v, the definition is given −∆[𝐴] by: 𝑣 = ∆𝑡. The reaction rate increases with: the concentration of the reactant; the temperature; the surface area of contact between reactants; and the concentration of any catalyst present. - Catalysts Catalysts are substances that increase the rate of a reaction without directly participating in it by lowering the activation energy. They do not change either ∆𝐻 or ∆𝐺, so they cannot make a non-exothermic reaction exothermic and spontaneous. Additionally, they do not alter the value of the equilibrium constant (Keq) of the reaction, thus they do not shift the equilibrium but rather accelerate its attainment. Catalysts contribute to making collisions more effective; they remain chemically unchanged at the end of the reaction; they do not appear in the overall reaction equations and are selective, meaning they only intervene in the reaction of interest. Biological reactions occur through specific catalysts called enzymes. 4. Chemical equilibrium Reactions that proceed until the complete exhaustion of the reactants are considered irreversible and are called completion reactions. However, many reactions are not complete because the products, once formed, can react with each other to reform the reactants. Reactions that occur in both directions are called reversible reactions and are indicated by two opposing arrows between the reactants and products. In these reactions, the forward and reverse reactions proceed at rates that eventually reach an equilibrium state. Once this state is reached, the concentrations of the reactants and products remain unchanged. In a reaction at equilibrium, the rates of the two reactions (forward and reverse) are equal. - Equilibrium Constant At a given temperature and at equilibrium, the ratio of the product of the molar concentrations of the products to that of the molar concentrations of the reactants, each raised to an exponent equal to its respective stoichiometric coefficient, is the thermodynamic equilibrium constant (Keq) of the reaction. α α [𝐶] ∗[𝐷] Mass action law: 𝐾𝑒𝑞 = α α ; 𝑝𝐾𝑒𝑞 =− 𝑙𝑜𝑔 𝐾𝑒𝑞. [𝐴] ∗[𝐵] The concentration values that satisfy this law are not the initial ones but those present when the system has reached equilibrium. - Properties of the equilibrium constant Keq is specific to each reaction and varies with temperature; The value of Keq indicates how far the reaction proceeds toward the formation of products: if it is very high, the reaction favours the products (shifted to the right), meaning at equilibrium it contains more products than reactants; if it is very low, it favours the reactants (shifted to the left), meaning at equilibrium it contains more reactants than products; if it is close to 1, the equilibrium contains approximately equal amounts of reactants and products. - Principle of mobile equilibrium The Chatelier's principle states that when a disturbance is applied to a system at equilibrium, the equilibrium conditions shift to counteract the disturbance; if one of the parameters of a system at equilibrium changes, it responds to counterbalance that change. The parameters that can vary include: the concentration of one of the components, temperature, and pressure. Adding a reactant or removing a product shifts the equilibrium to the right; removing a reactant or adding a product shifts the equilibrium to the left. In endothermic reactions, increasing temperature shifts the equilibrium to the right, increasing Keq; in exothermic reactions, increasing temperature shifts the equilibrium to the left, decreasing Keq. Acids and bases 1. Definitions There are various ways to define acids and bases. - Arrhenius definition Acid: a chemical species that dissociates in aqueous solution, releasing one or more hydrogen ions (H+), which actually exist as hydronium ions (H3O+). Base: a chemical species that releases one or more hydroxide ions (OH-) in aqueous solution. Acids and bases react with each other to form salts and water. This reaction is called neutralisation, and the energy released is termed the heat of neutralisation. The Arrhenius definition has some limitations: it is only applicable to aqueous solutions; it restricts the field of acidic and basic substances to those containing hydrogen atoms and dissociable OH groups; and it implies that acidity and basicity are absolute characteristics of various compounds. - Bronsted-Lowry definition Acid: a chemical species capable of transferring one or more H+ ions accepted by a base. Base: a chemical species capable of accepting one or more H+ ions donated by an acid. When an acid donates an H+ ion, it transforms into its conjugate base. When a base gains an H+ ion, it transforms into its conjugate acid. According to this theory, there are no standalone acids and bases, only acid/base conjugate pairs. Any reaction involving the transfer of H+ from an acid to a base is an acid-base reaction according to Bronsted-Lowry: acid + base → conjugate base + conjugate acid. An acid can act as a base under certain circumstances, and vice versa. - Amphoteric substances An amphoteric substance is a chemical species that behaves as a base in the presence of acids and as an acid in the presence of bases. For example, water behaves as a base with HNO3 and as an acid with NH3. - Lewis definition Acid: a chemical species that can accept a pair of non-bonding electrons. Base: a chemical species that can provide a pair of non-bonding electrons. Examples of Lewis acids include: H+ ion, certain metal cations, and molecules with an electron deficiency (e.g., BF3). An acid-base reaction according to Lewis involves the formation of a coordinate bond between a Lewis base (donor) and a Lewis acid (acceptor). 2. Strength of acids and bases Considering the compound XOH, the determining factor for dissociation is the electronegativity of X: - If X has low electronegativity, XOH behaves as a base; - If X has high electronegativity, XOH behaves as an acid; - If X has electronegativity similar to hydrogen, it is an amphoteric substance. An Arrhenius acid or base is stronger the more it dissociates in water. Strong acids and bases are substances that are fully dissociated into their constituent ions in aqueous solution. Weak acids and bases are substances that are only partially dissociated in aqueous solution. + − [𝐻 ]∗[𝐴 ] For a generic acid HA, the acid dissociation constant is given by: 𝐾𝑎 = [𝐻𝐴].​ The higher the value of Ka​, the stronger the acid: −4 𝐾𝑎 < 10 → weak acid −4 −1 10 < 𝐾𝑎 < 10 → medium acid −1 𝐾𝑎 > 10 → strong acid + − [𝑋 ]∗[𝑂𝐻 ] For a generic base XOH, the base dissociation constant is given by: 𝐾𝑏 = [𝑋𝑂𝐻]. The higher the value of KbK_bKb​, the stronger the base. 3. Ion product of water: Kw​ Pure water has very low electrical conductivity, indicating the presence of ions in solution. The presence of these ions is due to the fact that water, albeit very limited, dissociates according to the following autoionization reaction: + − −18 𝐻2𝑂 + 𝐻2𝑂 ↔ 𝐻3𝑂 + 𝑂𝐻. 𝐾𝑐 = 3. 25 ∗ 10 + − −14 The ion product of water is given by: 𝐾𝑤 = 𝐾𝑎 ∗ 𝐾𝑏 = [𝐻3𝑂 ] ∗ [𝑂𝐻 ] = 10 + − Its value increases with temperature. Based on the amounts of 𝐻3𝑂 and 𝑂𝐻 , solutions can be divided into three groups: + − Acidic solutions → [𝐻3𝑂 ] > [𝑂𝐻 ] + − Basic solutions → [𝐻3𝑂 ] < [𝑂𝐻 ] + − Neutral solutions → [𝐻3𝑂 ] = [𝑂𝐻 ]. 4. pH of an aqueous solution + The concentration of H+ ions, or H3O+, is often expressed as: 𝑝𝐻 =− 𝑙𝑜𝑔[𝐻3𝑂 ]. Based on the pH value, aqueous solutions are divided into three groups: Acidic solutions → pH7 - pH measurement pH metre: an instrument that measures pH by immersing an electrode in the solution being examined. Indicators: substances that change colour depending on the pH of the solution; they are weak acids or bases where the undissociated form has a different colour from the dissociated form. Indicator strips: strips of paper soaked in a mixture of indicators. The pH is determined by comparing the colour of a wet strip in the solution with a colour scale provided in the packaging. 5. Neutralisation and salt hydrolysis reactions An acid and a base react with each other to form salt and water in a neutralisation reaction. It is called this way because the H3O+ ions from the acidic solution combine with the OH- ions from the basic solution, resulting in a neutral solution (pH = 7). However, the solution is neutral only when equal amounts of a strong acid and a strong base react, or when equal amounts of a weak acid and a weak base react. When the reaction occurs between a weak acid and a strong base or between a strong acid and a weak base, the final solution has a pH different from This is due to the phenomenon of salt hydrolysis, where the ions derived from the dissociation of the salt react with water molecules, exchanging H3O+ ions. 6. Buffer Solutions A buffer solution consists of an aqueous solution of a weak acid and its salt with a strong base (or a weak base and its salt with a strong acid) in appropriate concentrations. These solutions have the property of maintaining their pH nearly unchanged even after the addition of small amounts of strong acids or bases. The pH also does not change if pure solvent is added to the solution: diluting a buffer solution keeps its pH constant. Organic Chemistry Organic chemistry is the chemistry of carbon compounds. Characteristics of carbon:  - Tendency to chain formation: it bonds with other carbon atoms to form open or closed chains. - External electronic configuration (2s² 2p²): belongs to group IV A and has 4 electrons in the outer shell. It tends to achieve an octet by forming covalent bonds. - Oxidation number: can vary from +4 to -4. - Hybridisation: sp³ (single bond), sp² (double bond), sp (triple bond). In organic compounds, carbon atoms are always hybridised and always form 4 covalent bonds. A carbon atom can be classified based on the number of carbon atoms it is bonded to: primary → bonded to one carbon atom secondary → bonded to two carbon atoms tertiary → bonded to three carbon atoms quaternary → bonded to four carbon atoms. 1. Nomenclature of organic compounds The systematic name (IUPAC nomenclature) of any organic compound consists of three parts: suffix → indicates the class of organic compounds to which the compound belongs; root → indicates the number of carbon atoms in the longest chain that can be identified in the compound; prefix → indicates the position of various substituents and/or functional groups in the main chain. 2. Functional groups A functional group is an atom or group of atoms that, attached to the carbon chain, is responsible for the chemical and physical properties of the compound. Functional group Name Class General formula Suffix C=C double bond Alkenes CnH2n -ene C≡C triple bond Alkynes CnH2n-2 -yne -X Halogen Alkyl halides CnH2n+1X -o -OH Hydroxyl group Alcohols R-OH -ol -O- Ether group Ethers R-O-R’ -ether H-C=O Carbonyl group Aldehydes R-CHO -al C=O Carbonyl group Ketones R-CO-R’ -one -COOH Carboxyl group Carboxylic acids R-COOH -oic acid -COOR’ Ester group Esters R-COOR’ -oate -CO-X Halogen Acyl halides R-CO-X -oyl halide -CO-NR’R’’ Amido group Amides R-CO-NR’R’’ -amide -NH2 Amino group Amines (primary) R-NH2 -amine C≡N Nitrile Nitriles R-C≡N -nitrile -NO2 Nitro groups Nitro derivatives R-NO2 -nitro -SH Thiol group Thiols R-SH -thiol 3. Reactivity of organic compounds The reactivity, meaning the tendency to transform, of any compound under the action of a reagent depends on the reaction conditions, the type of bond that must be broken, and the type of bond that must be formed. In most reactions, a positive or negative charge develops on the reacting molecule; if this charge is on a carbon atom, it is referred to as a carbocation or carbanion. These are formed when there is an electrolytic break of the covalent bond, in which one of the two atoms retains the electron pair and the other is left without. A carbon atom with an unpaired electron is a free radical, which can originate from a homolytic break of the covalent bond, where each of the two originally bonded atoms retains one of the two bonding electrons. In organic reactions, reagents are divided into two categories:  - Electrophilic reagents (Lewis acids): substances with an electron deficiency; a total or partial positive charge; the carbon atom in the carbonyl group -CO. They react with nucleophiles, which are rich in electrons. - Nucleophilic reagents (Lewis bases): substances capable of providing electrons because they possess a negative charge (total or partial); at least one lone pair; the nitrogen atom in the amine group -NH2. 4. Isomerism Isomers are two or more compounds that have the same molecular formula, and thus the same molecular weight, but different structural formulas or different spatial arrangements of atoms. - Conformational isomerism Conformational isomers differ in the relative spatial orientation of the groups of atoms; they can transform into one another by rotation around a single C-C bond without breaking any chemical bonds. Compounds containing double and triple bonds do not exhibit this type of isomerism. - Structural or chain isomerism Structural isomerism relates to the shape of the carbon chain, which can be linear, branched, or cyclic. Chain isomers are compounds that have the same molecular formula but different structural formulas. - Position isomerism Position isomers are compounds that have the same molecular formula but different positions in the chain of a substituent (atom or functional group) or a particular type of bond. - Stereoisomerism Stereoisomers are compounds in which the atoms are bonded in the same order or sequence but arranged differently in space. Unlike conformational isomers, they cannot be converted into one another without breaking chemical bonds. There are two types of stereoisomerism: ○ Optical isomerism (enantiomerism): Stereoisomers that are mirror images of each other but are not superimposable (like hands). The existence of enantiomers is linked to the presence of at least one chiral carbon atom (sp³ hybridised carbon bonded to four different atoms or groups). The enantiomer that rotates polarised light to the right is called dextrorotatory (+), while the one rotating to the left is called levorotatory (−). ○ Diastereoisomerism: Stereoisomers that are not mirror images of each other. A particular type is geometric isomerism (cis/trans): compounds that differ in the spatial arrangement of atoms or groups of atoms bonded to a cyclic hydrocarbon or to carbon atoms joined by a double bond. The cis isomer occurs when the two bulkier substituents are on the same side of the double bond; the trans isomer occurs when they are on opposite sides. Constitutional isomers: same molecular formula, but different arrangement of atoms: - Chain isomers - Position isomers - Functional group isomers Stereoisomers: same molecular formula, same bonds and functional groups, but different spatial arrangement: - Conformational isomers - Configurational isomers: ○ Optical isomers (enantiomers) ○ Diastereoisomers 5. Hydrocarbons Hydrocarbons are the simplest organic compounds. They are binary compounds made up solely of carbon and hydrogen and, based on their structure, are divided into different classes. An important initial distinction is made between aromatic hydrocarbons, which contain carbon rings with delocalized electron bonds, and aliphatic hydrocarbons, which do not contain delocalized electron bonds. Hydrocarbon Bond Hybridisation Geometry Bond angle Molecular orbitals Example Alkane single sp3 tetrahedral 109.5° all σ methane  Alkene double sp2 trigonal polar 120° 1σ+1𝜋 ethene Alkyne triple sp linear 180° 1σ+1𝜋 ethyne 6. Aliphatic hydrocarbons: alkanes, alkenes, alkynes According to IUPAC nomenclature, the name of a hydrocarbon is formed from a root that indicates the number of carbon atoms in the molecule, followed by a different suffix depending on the class. Prefixes: meth-, eth-, prop-, but-, pent-, hex-, hept-, oct-, non-, dec-. Suffixes: -ane, -ene, -yne. - Alkyl groups Substituents derived from hydrocarbons are called alkyl groups. The characteristic suffix of these groups is -yl, so the name of the group is obtained by replacing the suffix of the hydrocarbon with -yl: - Methyl: 𝐶𝐻3- - Ethyl: 𝐶𝐻3- 𝐶𝐻2- - N-Propyl: 𝐶𝐻3- 𝐶𝐻2- 𝐶𝐻2- - N-Butyl: 𝐶𝐻3- 𝐶𝐻2- 𝐶𝐻2- 𝐶𝐻2- - Nomenclature rules for organic compounds In the case of a branched hydrocarbon or when different functional groups are present, the compound is considered a derivative of the hydrocarbon corresponding to the longest carbon chain. The position of the substituents is indicated by the number of the carbon atom to which they are attached, after numbering the chain from the end closest to the substituent. If there are two or more identical substituents, their number is indicated with the prefixes di-, tri-, preceded by the numbers of the carbon atoms to which they are attached, separated by commas. The number must be repeated even if the identical substituents are attached to the same atom. - Alkanes Alkanes are saturated aliphatic hydrocarbons. They have the general formula 𝐶𝑛𝐻2𝑛+2; the characteristic suffix of the class is -ane; they contain only single covalent bonds C-C (σ bonds, sp³); they are devoid of functional groups. They may exhibit: conformational isomerism (free rotation around C-C bonds); structural isomerism (the chain can be linear, branched, open, closed, cyclic); optical isomerism (if it contains chiral carbon). The number of structural isomers increases with the number of carbon atoms in the molecule. This group includes methane (𝐶𝐻4), ethane (𝐶𝐻3- 𝐶𝐻3), propane (𝐶𝐻3- 𝐶𝐻2- 𝐶𝐻3), butane (𝐶𝐻3- 𝐶𝐻2- 𝐶𝐻2- 𝐶𝐻3) and pentane (𝐶𝐻3- 𝐶𝐻2- 𝐶𝐻2- 𝐶𝐻2- 𝐶𝐻3). - Cycloalkanes Cycloalkanes are the simplest cyclic aliphatic hydrocarbons. General formula: 𝐶𝑛𝐻2𝑛 with 𝑛 ≥ 3. The name is obtained by adding the prefix cyclo- to the name of the corresponding open-chain alkane. They can be represented by regular polygons and exhibit structural isomerism, conformational isomerism, cis/trans isomerism, and optical isomerism. - Alkenes Alkenes are unsaturated aliphatic hydrocarbons. They have the general formula 𝐶𝑛𝐻2𝑛 (𝑛 ≥ 2); the characteristic suffix is -ene; they contain one or more double bonds C=C (sp²; one σ bond and one 𝜋 bond). Nomenclature: the name is derived from that of the alkane with the same number of carbons by replacing the suffix -ane with -ene; the longest chain is numbered from the end closest to the double bond, and its position is indicated by the number of the carbon atom where the double bond begins. For example: 𝐶𝐻3- 𝐶𝐻2- 𝐶𝐻= 𝐶𝐻- 𝐶𝐻3 → 2-pentene. The simplest alkene is ethylene (𝐶𝐻2= 𝐶𝐻2), commonly known as ethylene. From butene onwards, alkenes exhibit structural isomerism, position isomerism, and cis/trans isomerism. Polyunsaturated alkenes contain more than one double bond; the simplest are dienes (𝐶𝑛𝐻2𝑛−2) with two double bonds. For example: 𝐶𝐻2=𝐶𝐻-𝐶𝐻=𝐶𝐻2 → 1,3-butadiene. - Alkynes Alkynes are unsaturated aliphatic hydrocarbons. They have the general formula 𝐶𝑛𝐻2𝑛−2 with 𝑛 ≥ 2; the characteristic suffix is -yne; they contain one or more triple bonds 𝐶≡𝐶 (one σ bond and two π bonds; sp). The simplest alkyne is ethyne (𝐶𝐻≡𝐶𝐻), also known as acetylene. Nomenclature: the name of the corresponding alkane is obtained by replacing the suffix -ane with -yne. From butyne onwards, alkynes exhibit position isomerism; starting from pentyne, they display structural isomerism. Example: 𝐶𝐻≡ 𝐶- 𝐶𝐻2- 𝐶𝐻3 → 1-butyne.  - Physical properties of aliphatic hydrocarbons Hydrocarbons are apolar molecular compounds, insoluble in water but soluble in apolar organic solvents (ether, benzene, chloroform). This is because hydrogen and carbon have similar and intermediate electronegativity values, making 𝐶-𝐶 and 𝐶-𝐻 bonds that are pure covalent bonds. Moreover, they are generally less dense than water; they have low melting and boiling points, although these increase with the number of carbons due to increased Van der Waals forces. At room temperature: 1-4 are gases; 5-16 are liquids; 17 and above are solids. - Chemical properties and reactions of alkanes Alkanes have very low chemical reactivity because 𝐶-𝐶 and 𝐶-𝐻 bonds are strong and weakly polarised. Due to this low reactivity, all reactions occur under drastic conditions: the presence of a catalyst and/or heat. Halogenation: a substitution reaction in which one or more hydrogen atoms are replaced by halogens (chlorine, bromine, or fluorine); the products are alkyl halides and the corresponding hydrogen halide. General reaction: 𝑅-𝐻 + 𝑋2(heat) → 𝑅-𝑋 + 𝐻𝑋. Combustion: an oxidation reaction where the hydrocarbon (fuel) reacts with oxygen (oxidant) to produce CO₂ and H₂O. The reaction is highly exothermic, but does not occur spontaneously at room temperature due to a high 1 activation energy (Ea). General reaction: 𝐶𝑛𝐻2𝑛+2 + (3𝑛 + 2 )𝑂2 → 𝑛𝐶𝑂2 + (𝑛 + 1)𝐻2𝑂 + ℎ𝑒𝑎𝑡. Dehydrogenation: in this reaction, hydrogen molecules are removed, producing unsaturated compounds (alkenes, alkynes). Example: 𝐶𝐻3-𝐶𝐻3 → 𝐻2 + 𝐶𝐻2=𝐶𝐻2. Cracking: a fragmentation reaction of a long hydrocarbon chain into smaller fragments, conducted at high temperature and possibly in the presence of catalysts. - Chemical properties and reactions of alkenes In addition to the reactions of alkanes, alkenes undergo specific reactions due to the presence of the double bond, which makes them more reactive. The characteristic reaction of alkenes is the addition reaction at the double bond. This has different names depending on the type of added reagent: Hydrogenation: alkene + 𝐻2 (catalyst) → corresponding alkane Halogenation: alkene + 𝑋2 → dihalide + Hydration: alkene + 𝐻- 𝑂𝐻 (𝐻 ) → alcohol Addition of hydrogen halides: alkene + 𝐻𝑋 → alkyl halide Polymerisation: 𝑅-𝐶𝐻2𝐶𝐻2 + 𝑛𝐶𝐻2=𝐶𝐻2 → 𝑅-(𝐶𝐻2𝐶𝐻2)𝑛-𝐶𝐻2𝐶𝐻2 A polymer is a large organic molecule made up of the repetition of a huge number of simpler units, called monomers. There are two types of polymers: addition polymers, formed through a series of chain addition reactions; condensation polymers, formed by a reaction between two different functional groups of two monomers with the elimination of a small molecule, generally water. Alkenes and alkynes are excellent raw materials for the synthesis of polymers. Natural polymers: starch, glycogen, cellulose, proteins, nucleic acids. Synthetic polymers: polyethylene (PE), polyvinyl chloride (PVC), polystyrene, polypropylene, Teflon, neoprene. - Chemical properties and reactions of alkynes The presence of the triple bond gives alkynes chemical properties similar to those of alkenes; they indeed undergo addition reactions. If one mole of alkyne is reacted with one mole of reagent, an unsaturated product (alkene) is obtained; if two moles are added, a saturated product (alkane) is obtained. Addition reactions: hydrogenation, halogenation, addition of hydrogen halides, hydration. 7. Aromatic hydrocarbons Aromatic compounds are special cyclic unsaturated hydrocarbons. The simplest is benzene, with the formula 𝐶6𝐻6: it is a cyclic compound with six carbon atoms in which the alternation of single and double bonds allows for the delocalization of 𝜋 bonding electrons throughout the molecule; for this reason, it is considered a resonance hybrid. It does not undergo spontaneous addition reactions but instead substitution reactions. - Derivatives of benzene Aromatic hydrocarbons derive from the substitution of one or more hydrogen atoms in the benzene ring with alkyl radicals or functional groups. Examples include toluene (-𝐶𝐻3), phenol (-𝑂𝐻), benzoic acid (-𝐶𝑂𝑂𝐻), styrene (-𝐶𝐻= 𝐶𝐻2), aniline (-𝑁𝐻2) and nitrobenzene (-𝑁𝑂2). When two substituents are present, their relative positions on the ring must be indicated. Two systems can be used: numbering the carbon atoms of the ring starting from one that has a substituent, or indicating the position of the second substituent relative to the first using the prefixes ortho- (1,2), meta- (1,3) and para- (1,4). Polycyclic aromatic compounds formed from two or more condensed or linked benzene rings can also be considered derivatives of benzene. Anthracene (3-), naphthalene (2-), phenanthrene (3), and benzo[a]pyrene (5). - Physical properties of aromatic hydrocarbons They are apolar compounds, and thus have physical properties similar to other hydrocarbons: insoluble in water and soluble in apolar organic solvents; less dense than water; melting and boiling points increase with the number of carbon atoms; at room temperature, benzene and its analogs are liquids, while polycyclic derivatives are generally solids. - Chemical properties of aromatic hydrocarbons The presence of the delocalized 𝜋 electron cloud makes benzene an electron-rich centre that behaves as a nucleophilic reagent (Lewis base) and tends to react with electrophilic reagents (Lewis acids). Benzene and its derivatives undergo electrophilic substitution reactions that allow for the maintenance of the aromatic sextet. This type of reaction has different names depending on the reagent involved: Nitration: reaction with nitric acid resulting in the substitution of a hydrogen atom with a 𝑁𝑂2 group: benzene + 𝐻𝑁𝑂3 (𝐻2𝑆𝑂4) → nitrobenzene + 𝐻2𝑂. Halogenation: reaction with a chloride leading to the substitution of a hydrogen atom in benzene with a chlorine atom: benzene + 𝐶𝑙2 (𝐴𝑙𝐶𝑙3) → 𝐶6𝐻5𝐶𝑙 + 𝐻𝐶𝑙. 8. Petroleum Petroleum is a mixture of aliphatic and aromatic hydrocarbons, both mono and polycyclic. The petroleum extracted from the ground is not usable as is and must undergo fractional distillation. 9. Derivatives of Hydrocarbons - Alcohols Alcohols have the general formula 𝑅-𝑂𝐻; the functional group is the hydroxyl group -𝑂𝐻; the characteristic suffix is -ol. Based on the type of carbon atom to which the -𝑂𝐻 group is attached, alcohols are classified as: Primary → -𝑂𝐻 group attached to a primary carbon atom; Secondary → -𝑂𝐻 group attached to a secondary carbon atom; Tertiary → -𝑂𝐻 group attached to a tertiary carbon atom. Nomenclature: the longest chain containing the -𝑂𝐻 group is numbered so that it receives the lowest possible number; the name of the alcohol is obtained by replacing the final letter of the corresponding hydrocarbon with the suffix -ol and indicating with a number the position of the --𝑂𝐻 group. In the case of polyfunctional alcohols, the suffixes -diol, -triol are used; for some alcohols, common names are employed. Physical properties: determined by the -𝑂𝐻 group, which makes the molecule polar and allows it to form hydrogen bonds; consequently, alcohols have higher melting and boiling points than hydrocarbons with the same number of carbon atoms, and alcohols with up to 3 carbon atoms are soluble in water, a characteristic that decreases with an increasing number of carbon atoms in the chain. Chemical properties: determined by the -𝑂𝐻 group; the presence of oxygen, which is much more electronegative than carbon, makes the molecule more reactive. Alcohols are amphoteric substances that can behave as weak acids or bases. The characteristic reaction is oxidation: primary alcohols can be oxidised to aldehydes, while secondary ones can be oxidised to ketones. Additionally, alcohols can be obtained by hydration of alkenes or by reduction of aldehydes and ketones. Common alcohols: methanol is used as a solvent, antifreeze, and as a raw material for the preparation of formaldehyde; it is toxic to the body. Ethanol can be obtained through biological fermentation or hydration of ethylene in acidic conditions. It is mainly used as fuel and in the food industry for the preparation of alcoholic beverages. Phenols: aromatic alcohols; general formula: 𝐴𝑟-𝑂𝐻; suffix -ol. The name derives from the corresponding hydrocarbons preceded by the prefix hydroxy- and one or more numbers identifying the position of -𝑂𝐻 groups. Thiol: general formula 𝑅-𝑆𝐻 (thioalcohols) or 𝐴𝑟-𝑆𝐻 (thiophenols); functional group is the thiol group -𝑆𝐻. The name is derived from the corresponding hydrocarbon by replacing the final letter with the suffix -thiol. - Ethers Ethers have the general formula 𝑅-𝑂-𝑅'. They are derived from water by the replacement of the two hydrogen atoms with two alkyl or aryl groups, which can be the same or different. The bond angle is approximately 110°. The name is derived from the word ether, preceded or followed by the names of the two radicals, listed in alphabetical order. For example: 𝐶𝐻3- 𝐶𝐻2-𝑂-𝐶𝐻2- 𝐶𝐻3 → diethyl ether. Ethers do not form hydrogen bonds, so they have lower melting and boiling points than alcohols. The ether bond is very strong; it requires treatment with strong acids at high temperatures to be cleaved. They are obtained through a condensation (dehydration) 𝐻2𝑆𝑂4 reaction between two alcohols: 𝑅 − 𝑂𝐻 + 𝑅'𝑂𝐻 ( 140° ) → 𝑅 − 𝑂 − 𝑅' + 𝐻2𝑂. - Aldehydes and ketones Aldehydes have the general formula R-CHO, while ketones have the general formula 𝑅 − 𝐶𝑂 − 𝑅'. The 𝐶=O group, common to both classes, is called the carbonyl group, which is strongly polarised, and the carbon is sp² hybridised. Nomenclature: the name of aldehydes is derived from the corresponding hydrocarbon by replacing the final letter with the suffix -al; the chain is numbered starting from the carbonyl carbon. Aromatic aldehydes and those of commercial use are often referred to by traditional names (propanal / propionic aldehyde). The name of ketones is derived from the corresponding hydrocarbon by replacing the final letter with the suffix -one; the chain is numbered from the end closest to the carbonyl carbon. Common names are still used: propanone / acetone / dimethyl ketone. Physical properties: the presence of the carbonyl group influences their properties. They have higher melting and boiling points than hydrocarbons with the same number of carbon atoms, but lower than the corresponding alcohols; being able to form hydrogen bonds with water, aldehydes and ketones with up to 5 carbon atoms are water-soluble, while those with more carbon atoms are soluble in organic solvents. Chemical properties: the presence of the highly electronegative oxygen atom means that the carbonyl group is strongly polarised and thus quite reactive. The typical reaction is nucleophilic addition, which occurs due to the attack by a Lewis base on the carbon of the carbonyl group, which has a partial positive charge and behaves as an electrophile. The addition of hydrogen, or reduction, of an aldehyde and a ketone provides a primary and a secondary alcohol, respectively. Aldehydes are easily oxidised to give the corresponding carboxylic acids, while ketones are oxidised only by strong oxidants. Common aldehydes and ketones: formaldehyde (methanal) is marketed as a 37% aqueous solution known as formalin and is used as a preservative and disinfectant. It is produced industrially by oxidation of methanol. Acetone (propanone) is used as a solvent because it can dissolve many organic substances, resins, and dyes. It is produced industrially by oxidation of isopropyl alcohol. - Carboxylic acids Carboxylic acids have the general formula 𝑅-𝐶𝑂𝑂𝐻. The functional group is the carboxyl group -𝐶𝑂𝑂𝐻, which is formed from a carbonyl group and a hydroxyl group. The carbon of the carboxyl group is sp² oxidised. Nomenclature: the name is derived from the name of the corresponding hydrocarbon preceded by the word acid, replacing the final letter with the suffix -oic; numbering of the chain starts from the carbon of the carboxyl group. Common names are also used; in this case, the position of substituents is specified using the Greek letters alpha, beta, gamma, starting from the carbon adjacent to the carboxyl group. Formic acid has antiseptic properties and is used as a preservative. Acetic acid is used as a raw material in numerous industrial syntheses. Citric acid is used in the pharmaceutical industry. Physical properties: polar molecules that can form hydrogen bonds. Solubility decreases with increasing carbon count; melting and boiling points are higher than those of corresponding hydrocarbons and alcohols. Chemical properties: they are weak acids that dissociate in aqueous solution: − + 𝑅𝐶𝑂𝑂𝐻 + 𝐻2𝑂 ↔ 𝑅𝐶𝑂𝑂 + 𝐻3𝑂. They can be obtained by oxidation of aldehydes. Treating an acid with a weak reducing agent can yield the corresponding aldehyde, while treating it with a strong reducer yields a primary alcohol. Fatty acids: These are carboxylic acids characterised by a long linear hydrocarbon chain. Saturated: palmitic (16), stearic (18). Unsaturated: oleic (18), linoleic (18), linolenic (18). - Esters Esters have the general formula 𝑅-𝐶𝑂𝑂-𝑅'. They are derived from acids by replacing the -𝑂𝐻 group with an -𝑂𝑅' group, known as the alkoxy group. The name is derived from the acid by replacing the suffix -ic with -ate, followed by the name of the alkoxy radical 𝑅'. For example: 𝐶𝐻3-𝐶𝑂𝑂- 𝐶𝐻2- 𝐶𝐻3 → ethyl acetate. Esters are obtained through a nucleophilic attack of an alcohol on a carboxylic acid, in the presence of an acid catalyst, in a reaction known as esterification. This is a condensation reaction between two compounds with the  + elimination of a water molecule: 𝑅-𝐶𝑂𝑂𝐻 + 𝑅'-𝑂𝐻 (𝐻 ) ↔ 𝑅-𝐶𝑂𝑂-𝑅' + 𝐻2𝑂. The reaction is reversible: the reverse of condensation is hydrolysis. Hydrolyzing an ester yields the starting acid and alcohol. Important are the esters of glycerol with fatty acids, known as triglycerides. - Anhydrides Anhydrides have the general formula 𝑅-𝐶𝑂-𝑂-𝐶𝑂-𝑅'. The name is derived from that of the corresponding acid by replacing the word acid with anhydride. For example: 𝐶𝐻3-𝐶𝑂-𝑂-𝐶𝑂-𝐶𝐻3 → acetic anhydride. Anhydrides are prepared through the condensation of carboxylic acids, according to the reaction: + 𝐻 𝑅-𝐶𝑂𝑂𝐻 + 𝑅'𝐶𝑂𝑂𝐻 ( ℎ𝑒𝑎𝑡 ) ↔ 𝑅-𝐶𝑂-𝑂-𝐶𝑂-𝑅' + 𝐻2𝑂. Hydrolyzing anhydrides yields the corresponding acids. - Amines Amines are organic derivatives of ammonia (𝑁𝐻3) obtained by replacing one or more hydrogen atoms with corresponding alkyl or aryl groups. There are three groups of amines: Primary → 𝑅-𝑁𝐻2 Secondary → 𝑅-𝑁𝐻-𝑅' Tertiary → 𝑅-𝑁-𝑅'𝑅'' The most important amine is aniline, used in the preparation of dyes, varnishes, and paints. Many biologically important compounds are amines: serotonin and morphine. Amines are polar molecules that can form hydrogen bonds with themselves, except for tertiary amines, and with water. However, the hydrogen bond is weaker than in alcohols because nitrogen is less electronegative than oxygen. As a result, they have higher melting and boiling points than hydrocarbons of equal molecular weight, although lower than those of the corresponding alcohols; they are water-soluble when the carbon atoms are more than 6. The lone pair of electrons on nitrogen explains the basic properties of amines, which can accept H⁺ ions from + − water according to the reaction: 𝑅-𝑁𝐻2 + 𝐻2𝑂 ↔ 𝑅-𝑁𝐻3 + 𝑂𝐻. - Amides Amides have the general formula 𝑅-𝐶𝑂-𝑁𝑅'𝑅''. They are derived from acids by replacing the -𝑂𝐻 group of a carboxylic acid with an amine group -𝑁𝐻2, with -𝑁𝐻𝑅 (primary amine), or with -𝑁𝑅'𝑅'' (secondary). The name is derived by replacing the suffix -oic of the acid with -amide. Amides are obtained from a condensation reaction between a carboxylic acid or its derivative and ammonia or a primary/secondary amine. The reaction is reversible: they are hydrolyzed in a basic or acidic environment, releasing the corresponding carboxylic acid and ammonia or ammonium ions, respectively. - Nitriles (cyanides) Nitriles have the general formula 𝑅-𝐶≡𝑁 and are derived from hydrocyanic acid (HCN) by replacing the hydrogen atom with an alkyl or aryl group. The functional group is 𝐶≡𝑁, known as cyano or nitrile. Suffix: -nitrile. - Heterocyclic compounds Heterocycles are cyclic organic compounds, either aliphatic or aromatic, in which the ring contains one or more atoms different from carbon, typically nitrogen, oxygen, or sulphur. The most common heterocycles are five-membered and six-membered. Based on the structure, they are divided into non-aromatic and aromatic. They generally undergo electrophilic substitution reactions. In addition to heterocycles formed from a single ring (pyrimidines), there are those formed from two or more rings (purines). Organic compounds of biological interest The most important compounds from a biological perspective, also known as biomolecules, belong to four classes: carbohydrates, proteins, nucleic acids and lipids. The compounds in the first three classes are polymers formed from simpler units (monosaccharides, amino acids, nucleotides). The class of lipids includes compounds with heterogeneous chemical structures, characterised by similar physical properties. 1. Carbohydrates Carbohydrates, also known as saccharides, glucides, or sugars, are ternary compounds containing carbon, oxygen, and hydrogen. General formula 𝐶𝑛(𝐻2𝑂) where 𝑚 ≤ 𝑛. They contain various alcohol groups and an