Classification of Elements and the Periodic Table
The periodic table is one of the most significant achievements in chemistry, organizing all known elements in a coherent and predictive manner. Understanding its structure and the principles behind it is crucial for comprehending chemical behavior.
4.1 Modern Periodic Law and Modern Periodic Table
The Modern Periodic Law states that the physical and chemical properties of elements are periodic functions of their atomic numbers. This law, proposed by Moseley, rectified the anomalies of Mendeleev's table by arranging elements based on increasing atomic number rather than atomic mass.
4.1.1 Classification of elements into different groups, periods and blocks
- Periods: Horizontal rows (1-7) representing the principal quantum number (n) of the outermost electron shell.
- Groups: Vertical columns (1-18) where elements share similar chemical properties due to having the same number of valence electrons.
- Blocks: Elements are categorized into s, p, d, and f blocks based on the orbital in which the last electron enters.
- s-block: Groups 1 and 2 (Alkali and Alkaline Earth Metals). Highly reactive metals.
- p-block: Groups 13 to 18 (Boron to Noble Gases). Contains metals, non-metals, and metalloids.
- d-block: Groups 3 to 12 (Transition Metals). Exhibit variable oxidation states and form colored compounds.
- f-block: Lanthanides and Actinides (Inner Transition Metals). Often found below the main body of the periodic table.
4.2 IUPAC Classification of Elements
The International Union of Pure and Applied Chemistry (IUPAC) recommends a system where groups are numbered from 1 to 18. This modern system provides a clear and unambiguous way to refer to groups, replacing older notations (e.g., IA, IIA, IIIB).
4.3 Nuclear Charge and Effective Nuclear Charge
- Nuclear Charge (Z): The total positive charge in the nucleus, equal to the atomic number. It attracts electrons towards the nucleus.
- Effective Nuclear Charge (Zeff): The net positive charge experienced by an electron in a multi-electron atom. It is less than the actual nuclear charge due to shielding by inner electrons. Zeff = Z - S (where S is the shielding constant).
4.4 Periodic Trend and Periodicity
Periodicity refers to the recurring trends in the physical and chemical properties of elements when arranged by atomic number.
4.4.1 Atomic Radii
The distance from the center of the nucleus to the outermost electron shell. Generally, atomic radius:
- Decreases across a period: Due to increasing Zeff pulling the electron cloud closer to the nucleus.
- Increases down a group: Due to the addition of new electron shells.
4.4.2 Ionic Radii
The radius of an ion. It differs from atomic radius because of the gain or loss of electrons.
- Cations (positive ions): Smaller than their parent atoms because they lose valence electrons, reducing electron-electron repulsion and increasing Zeff per electron.
- Anions (negative ions): Larger than their parent atoms because they gain electrons, increasing electron-electron repulsion and decreasing Zeff per electron.
4.4.3 Ionization Energy (IE)
The minimum energy required to remove one electron from a gaseous atom in its ground state to form a gaseous ion.
- Increases across a period: Due to increasing Zeff and decreasing atomic size, making it harder to remove an electron.
- Decreases down a group: Due to increasing atomic size and greater shielding, making it easier to remove an electron.
- Exceptions: Anomalies occur due to stable half-filled or fully-filled subshells (e.g., IE of N > IE of O, IE of Be > IE of B).
4.4.4 Electron Affinity (EA)
The energy change when an electron is added to a neutral gaseous atom to form a negative ion. A more negative (exothermic) EA indicates a greater tendency to accept an electron.
- Generally becomes more negative across a period: Due to increasing Zeff.
- Generally becomes less negative down a group: Due to increasing atomic size and weaker attraction for the incoming electron.
- Halogens have very high (negative) EA. Noble gases have positive EA.
4.4.5 Electronegativity (EN)
The ability of an atom in a molecule to attract shared electrons towards itself. It's a relative measure.
- Increases across a period: Due to increasing Zeff.
- Decreases down a group: Due to increasing atomic size and shielding.
- Fluorine (F) is the most electronegative element.
4.4.6 Metallic Characters (General trend and explanation only)
Metallic character is associated with the ease of losing electrons (low IE, low EN). Metals are typically shiny, malleable, ductile, and good conductors of heat and electricity.
- Decreases across a period: Elements become more non-metallic (tendency to gain electrons).
- Increases down a group: Elements become more metallic (tendency to lose electrons more easily).
Chemical Bonding and Shapes of Molecules
Chemical bonds are the forces that hold atoms together to form molecules and compounds. Understanding these bonds is fundamental to predicting the properties and reactions of substances.
5.1 Valence shell, valence electron and octet theory
- Valence Shell: The outermost electron shell of an atom.
- Valence Electrons: Electrons in the valence shell, primarily involved in chemical bonding.
- Octet Theory (or Octet Rule): Atoms tend to gain, lose, or share electrons in order to achieve a full outer electron shell with eight electrons (like noble gases), which is a state of maximum stability. (Exceptions exist, especially for elements in period 3 and beyond, and for elements like H, He, Li, Be, B).
5.2 Ionic bond and its properties
An ionic bond forms between a metal and a non-metal, involving the complete transfer of one or more electrons from the metal to the non-metal, resulting in the formation of oppositely charged ions which are held together by electrostatic forces.
- Properties of Ionic Compounds: Usually solid at room temperature, high melting and boiling points, good conductors of electricity when molten or dissolved in water, often soluble in polar solvents, and brittle.
5.3 Covalent bond and coordinate covalent bond
- Covalent Bond: Formed by the sharing of one or more pairs of electrons between two non-metal atoms. This sharing allows both atoms to achieve a stable electron configuration. Can be single, double, or triple bonds.
- Coordinate Covalent Bond (Dative Bond): A type of covalent bond where both shared electrons come from only one of the participating atoms. Once formed, it is indistinguishable from a regular covalent bond (e.g., in ammonium ion, NH4+).
5.4 Properties of covalent compounds
- Typically gases, liquids, or low melting point solids at room temperature.
- Lower melting and boiling points compared to ionic compounds.
- Poor conductors of electricity in all states.
- Often insoluble in polar solvents but soluble in non-polar solvents.
5.5 Lewis dot structure of some common compounds of s and p block elements
Lewis dot structures represent valence electrons as dots around the atomic symbol, showing how electrons are shared or transferred to achieve octets.
Examples:
- H2O: Oxygen in the center, bonded to two hydrogens, with two lone pairs on oxygen.
- CO2: Carbon in the center, double-bonded to two oxygens, with two lone pairs on each oxygen.
- NH3: Nitrogen in the center, bonded to three hydrogens, with one lone pair on nitrogen.
- CH4: Carbon in the center, bonded to four hydrogens.
5.6 Resonance
Resonance occurs when a single Lewis structure cannot adequately describe the bonding in a molecule or ion. Instead, multiple valid Lewis structures (resonance structures) can be drawn, and the actual molecule is a hybrid of these structures. This phenomenon leads to delocalized electrons and increased stability (e.g., carbonate ion CO32-, benzene C6H6).
5.7 VSEPR theory and shapes of some simple molecules
Valence Shell Electron Pair Repulsion (VSEPR) Theory predicts the geometry of molecules based on the repulsion between electron pairs (both bonding and non-bonding) in the valence shell of the central atom. Electron pairs will arrange themselves as far apart as possible to minimize repulsion.
Examples of Molecular Shapes:
- BeF2: Linear (2 electron groups, 0 lone pairs, 180° bond angle)
- BF3: Trigonal Planar (3 electron groups, 0 lone pairs, 120° bond angle)
- CH4: Tetrahedral (4 electron groups, 0 lone pairs, 109.5° bond angle)
- CH3Cl: Tetrahedral (slight distortion due to different atom sizes)
- PCl5: Trigonal Bipyramidal (5 electron groups, 0 lone pairs)
- SF6: Octahedral (6 electron groups, 0 lone pairs)
- H2O: Bent (4 electron groups, 2 lone pairs, ~104.5° bond angle)
- NH3: Trigonal Pyramidal (4 electron groups, 1 lone pair, ~107° bond angle)
- CO2: Linear (2 electron groups, 0 lone pairs, 180° bond angle)
- H2S: Bent (similar to H2O)
- PH3: Trigonal Pyramidal (similar to NH3)
5.8 Elementary idea of Valence Bond Theory
Valence Bond Theory (VBT) describes covalent bonds as the overlap of atomic orbitals (half-filled) from different atoms. The greater the overlap, the stronger the bond. It introduces the concept of sigma (σ) and pi (π) bonds.
5.9 Hybridization involving s and p orbitals only
Hybridization is the concept of mixing atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons to form chemical bonds. This helps explain observed molecular geometries that cannot be explained by simple atomic orbital overlap.
Types involving s and p orbitals:
- sp hybridization: One s orbital and one p orbital mix to form two sp hybrid orbitals (linear geometry, e.g., BeCl2, C2H2).
- sp2 hybridization: One s orbital and two p orbitals mix to form three sp2 hybrid orbitals (trigonal planar geometry, e.g., BF3, C2H4).
- sp3 hybridization: One s orbital and three p orbitals mix to form four sp3 hybrid orbitals (tetrahedral geometry, e.g., CH4, NH3, H2O).
5.10 Bond characteristics:
5.10.1 Bond length
The average distance between the nuclei of two bonded atoms. Depends on atomic radii, bond order (single > double > triple), and hybridization.
5.10.2 Ionic character
The degree to which a covalent bond has ionic properties. It increases with the difference in electronegativity between the two bonded atoms. A bond with a 0% EN difference is purely covalent, while a very large difference implies significant ionic character.
5.10.3 Dipole moment
A measure of the polarity of a molecule, resulting from the uneven distribution of electron density. It is a vector quantity. A molecule with a net dipole moment is considered polar (e.g., H2O, NH3), while symmetrical molecules with polar bonds might have a zero net dipole moment (e.g., CO2, CCl4).
5.11 Vander Waal’s force and molecular solids
Van der Waals forces are weak intermolecular forces (IMFs) that include dipole-dipole forces, London dispersion forces (present in all molecules, strongest in large nonpolar molecules), and dipole-induced dipole forces.
- Molecular Solids: Solids held together by intermolecular forces. They typically have low melting points, are soft, and are poor conductors of electricity (e.g., ice, sugar, dry ice).
5.12 Hydrogen bonding and its application
Hydrogen bonding is a special, strong type of dipole-dipole interaction that occurs when hydrogen is bonded to a highly electronegative atom (N, O, or F). The H atom becomes highly positive and is attracted to a lone pair on another N, O, or F atom.
Applications/Consequences: High boiling point of water, structure of DNA, protein folding, solubility of alcohols in water.
5.13 Metallic bonding and properties of metallic solids
Metallic bonding is described by the