curs chimie an I ISE, EI, IEI, TTC, ISER, ETI

1. INTRODUCTION TO ATOMIC STRUCTURE

1.1 Atomic Structure

• J. Dalton's theory considers atoms as unchangeable and indestructible particles in chemical reactions.

• Discoveries from the late 19th and early 20th centuries (cathode rays, atomic spectra, X-ray spectra) showed the complexity of atomic structure, the discrete nature of energy exchanges with the environment, and the corpuscular-wave nature of elementary particles.

• Elementary particles consist of:

  • Proton (p+), mass = 1.0073 u.a.m., charge = +1.

  • Neutron (n), mass = 1.0087 u.a.m., no charge.

  • Electron (e-), negligible mass (0.00055 u.a.m.), charge = -1.

• Protons and neutrons form the atomic nucleus; electrons form the electronic shell.

• Rutherford's experiments revealed the atom contains a small, massive, positively charged nucleus, surrounded by electrons.

• The atomic number (Z) represents the number of protons in the nucleus, the number of electrons in the electronic shell, and the element's position in the periodic table.

• Atoms with the same atomic number (Z) form a chemical element.

• The mass number (A) is the sum of protons (Z) and neutrons (N) in a nucleus: A = Z + N.

• Atoms with the same atomic number but different mass numbers are called isotopes.

• Isotopes have similar physical and chemical properties.

• Important applications of isotopes include:

  • Tracers in industries, agriculture, medicine.

  • Age determination of carbon-containing materials (14C) in archaeology and paleontology.

  • Energy production and medical applications (radioactive isotopes).

• Atomic mass is calculated as the weighted average of the atomic masses of an element's isotopes.

• The Bohr atomic model describes hydrogen's atom, assuming the electron moves in a circular orbit around the nucleus with quantized energy levels determined by quantum number (n).

• Energy of an electron in a hydrogen-like system is given by the formula.

1.2 Nuclear Composition of the First Ten Elements

*The 12th part of the isotope's mass represents the atomic mass unit (u.a.m).

• Bohr postulated that an atom emits energy in the form of a photon when an electron transitions from a higher to a lower energy level.

• The theory was widely accepted after spectral lines obtained for hydrogen validated Bohr's theory.

• Light and other forms of electromagnetic radiation travel in a vacuum with a speed c = 2.998 x 108 m/s.

• Radiation behaves like a wave characterized by frequency (ν) and wavelength (λ), related by c = λν.

• Electromagnetic radiation also possesses particle properties (photons).

• The energy of a photon is correlated with the frequency of the radiation: E = hν, where h is Planck's constant.

• The hydrogen line spectrum is obtained by passing light produced in a discharge tube through a prism; measuring the wavelength corresponding to each line can determine the energy of the respective photon.

2. KINETIC AND QUANTUM MODELS

2.1 Quantum Model

• Atoms with two or more electrons cannot be adequately described by Bohr's model; thus, a quantum model was adopted wherein an electron's behavior is described by a wave function (ψ) characterizing an orbital, or the region in space occupied by the electron.

• The wave function allows calculation of an electron's energy and probability of finding the electron in a specific position in the orbital.

• The electronic distribution in an orbital is described as electronic density.

• An orbital is characterized by four quantum numbers:

  • Principal quantum number n corresponds to the number of layers and the period in which the element is found.

  • Secondary quantum number l (ranging from 0 to n-1) corresponds to the type of orbital.

  • Magnetic quantum number m (ranging from -l to +l) describes the orbital orientation.

  • Spin quantum number s (s = 1/2) describes the electron's movement around the axis.

• Orbitals can have various shapes: spherical (s), bilobed (p), four-lobed (d), and eight-lobed (f).

• The energy of electrons in various orbitals increases in the order: s electrons < p electrons < d electrons < f electrons.

• Orbitals of the same type form a subshell: s type (with one orbital), p type (with three orbitals), d type (with five orbitals), and f type (with seven orbitals).

• The electronic distribution of an atom in shells and subshells constitutes the electronic configuration.

• The number of electrons in a shell is equal to 2n², where n is the shell number.

• Two rules govern determining the electronic configuration:

  • Pauli's Exclusion Principle: only two electrons with opposite spin can occupy an orbital.

  • Hund's Rule: all orbitals of a subshell are first occupied with one electron each before pairing with opposite spin electrons.

2.2 Periodic System of Elements

• D. Mendeleev and L. Meyer observed that the properties of elements are periodic functions of atomic mass.

• The modern statement of the law of periodicity is that the properties of elements are periodic functions of the atomic number (Z) and depend on the electronic configuration of atoms.

• The periodic table is an arrangement of elements in ascending order of atomic numbers, positioning similar elements in vertical columns.

• The table consists of seven horizontal rows (periods or series) and 32 vertical columns (groups).

• The first short period contains hydrogen (H) and helium (He). Each of the second and third periods contains eight elements. The fourth and fifth periods each consist of 18 elements, including the first and second series of transitional elements. The sixth period contains 32 elements, including the third transitional series and the first series of inner transitional elements: lanthanides or rare earths. The seventh period and final one contains the fourth transitional series and the second series of inner transitional elements: actinides.

• Elements can be classified into four blocks based on the type of orbital containing the distinguishing electron:

  • Noble gases block has fully occupied outer orbitals (ns2p6).

  • s and p block elements include main group elements (IA–VIIIA).

  • d block elements consist of transitional elements from secondary groups (IB-VIIIB).

  • f block elements include inner transitional elements: lanthanides and actinides.

• Observing the variation of physical and chemical properties of elements concerning their position in the periodic system reveals two representative groups: properties determined by the nucleus vary linearly, aperiodically, while those determined by the outer electronic configuration of atoms vary periodically.

2.2.1 Non-Periodic Properties of Elements

• The atomic number (Z) represents the fundamental property of atoms and increases monotonically by one from one element to another.

• The atomic mass (A) increases monotonically from the first element to the last.

• Initially considered a fundamental property, it was later established as a derived property depending on the isotopic composition of elements.

• X-ray spectra have shown a linear dependency of emitted lines' frequency for elements with Z > 11 concerning nuclear charge (Z).

2.2.2 Periodic Properties of Elements

• This category includes various physical properties: atomic radius, atomic volume, ionic radius, ionization energy, electron affinity, redox potential, optical spectra, magnetic properties, and chemical properties (combination ability, electrochemical character, oxidation-reduction and acid-base character).

• The atomic radius decreases in a period with increasing Z and increases in main groups with Z due to an increase in the number of electronic levels, while in secondary groups, atomic radii vary chaotically.

2.2.3 Trends in Ionization Energy and Electron Affinity

• Ionization energy is the energy required to remove one or more electrons from a neutral atom. It generally increases across a period and decreases down a group in the periodic table.

• Electron affinity is the energy released or absorbed when an atom accepts an electron in its valence shell, forming an anion.

2.2.3 Trends in Redox Properties

• The values of ionization energy and electron affinity allow the determination of the electrochemical character (electropositive or electronegative) of elements.

• Electropositive character is the tendency of an isolated atom to give away electrons, increasing across a period from group VII to group I, and down a group with increasing Z.

• Elements exhibiting electropositive character are metals, while electronegative character is characteristic of non-metals.

2.2.4 Valence and Oxidation States

• Valence is the capacity of an element to combine with another element; the oxidation state (N.O.) represents the actual or formal electric charge attributed to an atom in a compound based on established rules:

  • elements in their elemental state have N.O. = 0; hydrogen in ionic hydrides has −1 and in covalent hydrides has +1; oxygen has N.O. = −2; metals in main groups have N.O. equal to their ion charge, while non-metals have N.O. equal to 8 minus their group number.

• Maximum positive N.O. equals the group number of the element.

3. CHEMICAL BONDS

3.1 Chemical Bonds Theory

• Chemical bonds, ionic and covalent, are formed by atoms striving to achieve a stable octet configuration in their outer electronic shells.

3.1.1 Ionic Bonding

• An ionic bond forms through the transfer of electrons between atoms, as in the case of sodium chloride.

• The number of electrons transferred by an atom represents the electrovalence and dictates the charge on each ion.

• A stable octet configuration applies only to main groups; for transitional elements, stable configurations include d5, f7, f14, and hydrogen achieves a stable configuration of a duet.

• Ionic bond charactistics include: electrostatic attraction between ions, isotropy, strong nature over interionic distances.

3.1.2 Covalent Bonding

• A covalent bond is formed by sharing electrons between atoms to establish a stable configuration.

• The electronic theory establishes three types of covalency:

  • Non-polar covalent bonds occur between atoms of the same kind with uniform electron distribution between them.

  • Polar covalent bonds occur between different chemical nature atoms, leading to partial negative and positive charges.

  • Coordinative covalent bonds occur when one atom provides a pair of unshared electrons to an electron-deficient atom.

3.1.3 Characteristics of Covalent Bonds

• Covalent bonds exhibit rigidity, spatial orientation, saturation, and strength, and their breakage leads to the chemical transformation of the substance.

3.2 Metallic Bonding

• Metallic bonding characteristic of metals has been explained by three theories:

  • Also, the gas electron theory considers metals as crystalline networks composed of positively charged ions and freely moving valence electrons.

  • The valence bond theory regards metallic bonding as a delocalized covalence in the directions where atoms reside in the crystalline network.

  • The molecular orbitals method or band theory views metallic bonding as a strong delocalized covalent bond formed within the field of all nuclei, with bonding molecular orbitals establishing the valence band (banded conductor).

4. INTERMOLECULAR BONDS AND STATES OF MATTER

4.1 Gaseous State

• The gaseous state is typical for monoatomic gases (rare gases) and covalent substances with small molecular mass: H2, O2, N2, Cl2, HCl, NH3, and some hydrocarbons.

• At high temperatures, most substances transition into a homogeneous gaseous state governed by laws not dependent on the chemical nature of the gas.

• Specific properties of the gaseous state, dictated by translational, vibrational, and rotational movements of molecules with weak Van der Waals forces, include: diffusion, expansibility, compressibility, and transparency.

• Gases have no fixed shape or volume; they occupy all available space, adapting to its shape.

4.1.1 Pressure Measurement and The Ideal Gas Law

• The pressure of a gas is expressed in atmospheres, torr, bar, Pascals, or mm of mercury column: 1 atm = 760 mm Hg = 760 torr = 101.325 kPa = 1 bar.

• Under the same conditions of pressure and temperature, equal volumes of gases contain the same number of molecules (Avogadro's law).

• The ideal gas state is a limiting state derived from the kinetic-molecular theory, considering gas molecules as negligible volume points without cohesive forces.

• Gas molecules move randomly, changing direction exclusively due to collisions with themselves or container walls, with elastic collisions.

• The average kinetic energy of gas molecules is directly proportional to the temperature at which the gas is found.

4.1.2 Liquid State

• The liquid state is intermediate between a gaseous and a solid state. Particles have less freedom, their movements are mainly vibrational and translational or sliding.

• Liquids do not have a fixed shape but maintain a definite volume.

• Depending on the nature of the particles, liquids can be ionic, atomic, metallic, or polar/non-polar molecular.

• Microscopic ordered structures called micro-clusters occur at close distances in liquids.

4.1.3 Solid State

• The solid state is the most condensed state of matter due to strong cohesive forces, with the only possible movements of particles being vibrational.

• Solids can be classified into two forms based on the order of particles: crystalline and amorphous.

• Crystals have fixed melting points, transition sharply to liquid, and break along flat surfaces (cleavage), exhibiting anisotropy.

5. THERMOCHEMISTRY

5.1 First Law of Thermodynamics

• Chemical thermodynamics studies energy transformations accompanying chemical reactions, the stability conditions of chemical systems.

• The first principle of thermodynamics states: energy cannot be lost; it transforms, and the total energy remains constant.

5.2 Thermochemical Equations

• Thermochemical equations indicate the thermal effect (expressed in Joules or calories) and the state of reactants/products:

  • combustion of carbon: C(s) + O2(g) = CO2(g)

  • reaction of sodium with water: 2Na(s) + 2H2O(l) = 2NaOH(s) + H2(g)

  • decomposition of limestone: CaCO3(s) → CaO(s) + CO2(g)

5.3 Heat of Reaction

• Heat required for the formation of one mole of substance is called heat of formation.

• Standard enthalpy of formation is noted as H.

• Calorimetric experiments occur in solution at constant pressure, as reactions occur without significant volume changes.

6. CORROSION

6.1 Corrosion Processes

• Corrosion is a destructive process involving the transformation from a metal to its oxidized form in interaction with the surrounding environment.

• Two principal procedures describe the corrosion process: chemical and electrochemical corrosion.

• Factors influencing corrosion include material purity, environmental pH, temperature, and presence of corrosive agents.

6.2 Protection Against Corrosion

• Use of coatings, choice of materials and treatments can significantly reduce corrosion:

  • Metallic Coatings: including galvanization (Zn), engine oiling, and protective coatings.

  • Non-Metallic Coatings: such as varnishes, plastics, and bituminous materials to prevent corrosion.

  • Electrochemical Protection: polarization anodic or cathodic. The choice of protection method depends on maximizing efficiency during design.