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• Atomic Theory Timeline


  • 400 BC - Democritus: Atomos - Building blocks of matter, The shape of an atom explain the elements behavior
  • 1803 - John Dalton: Solid Sphere - Atom is solid indivisible sphere
  • 1897 - J.J Thomson: Plum Pudding - Negative electrons are embedded in a sea of positive charges
  • 1911 - Ernest Rutherford: Nuclear - Positive charges are located in central nucleus
  • 1913 - Niels Bohr: Planetary - Electrons having circular orbits with different energy level
  • 1926 - Erwin Schrodinger: Quantum - Electrons are in the clouds, Less dense

  
• Atomic Orbitals


  • Electrons act like waves
  • Exact location of electrons cannot be determined

  
• Orbitals


  • Region in an atom where there is a high probability of finding electrons
  • Atomic Orbitals (s, p, d, or f) - "f" most diffused shape

  
• Electron Configuration Notation


  Describes how electrons are distributed in atomic orbitals

  
• Principal Quantum Number


  Indicates the energy level or shell where an atomic orbital can be found
• Azimuthal Quantum Number


  • Specifies the sublevel (or subshell) within a particular principal energy level
  • 0 = s
  • 1 = p
  • 2 = d
  • 3 = f
• Magnetic Quantum Number


  • Indicates the specific orbital within the sublevel where the electron is found
  • n-1
  • s= 0
  • p= -1, 0, 1
  • d= -2, -1, 0, 1, 2
  • f= -3, -2, -1, 0, 1, 2, 3 (represented as box)
• Spin Quantum number


  • Pauli exclusion principle, only a maximum of two electrons can occupy an orbital, and they must have opposite spins to minimize repulsion between them
  • ⬆⬇
• Rules that determine how electrons fill atomic orbitals

  • Aufbau Principle - Electrons fill the lowest energy orbitals first
  • Hund's Rule - Atomic orbitals maximize the number of electrons with the same spin, Electrons will singly occupy each orbital and with parallel spins before they pair up
  • Pauli Exclusion Principle - Electrons will pair with opposite spins

  
• Octet Rule


  • Atoms combine to form compounds and follow the noble gas configuration, a stable configuration
  • There are two ways of attaining such stable configuration: Transfer of electron/s (gaining or losing), Sharing of electrons
  • Valence Electrons - the electrons involved in chemical bonding
• Lewis Structure


  • A dot is placed in each of the four sides of the element symbol before pairing with another as needed to represent all the valence electrons of the elements
  • There are no strict rules on which sides to pair up first so the dot symbol for oxygen may be written in several equivalent forms

  
• Lewis Electron Dot Structure (LED)

  • Used to illustrate how valence electrons participate in chemical bonding
  • The core/central element is represented by its chemical symbol
  • The valence electrons are represented by dots around the chemical symbol

  
• Stability of Noble Gases

  • Elements in Group 8A also called as noble gases
  • Most stable elements in the periodic table
  • Non Reactive under ordinary conditions; hence, the description "inert"
  • Except for helium with only two electrons in its orbital
  • Noble gases with a general valence configuration of ns2np2 have 8 valence electrons
  • Such octet configuration is the most stable arrangement an atom can have
  • The atoms of the other elements in the periodic table tend to achieve the configuration of the nearest noble gas by reacting with the same element or with other elements to form new compounds
  • This principle is referred to as the octet rule
• Gilbert Lewis


  • American chemist
  • In 1916, he developed a system of representing the valence electrons of atoms using diagrams called Lewis electrons–dot structures or Lewis structures

  
• Lewis structure


  • Consists of a symbol of an element surrounded by one or more dots: each dot corresponds to a valence electron in an atom of an element
  • Only two dots are placed on each of the four sides

  
• Metals


  Have one to three valence electron, which can be easily removed because of their relatively low ionization energy
• Nonmetals


  Having high electron affinity can gain valence electrons to fill their s and p orbitals and form an octet
• Formation of Ionic Compound

  Electrons given off by a metal atom are gained by the nonmetal, forming ionic bond in the process

  
• Ionic compounds and Lewis Structures

  Ionic generally exists between a metal and a nonmetal as a result of their high electronegativity difference

  
• Properties of Ionic Compound

  • High Melting Point and Boiling Point - The ionic bonds that bind the ions are strong such that high energy is required to separate the ions and allow them to move freely and form a liquid and gas
  • Conducts Electricity - Solid ionic compounds generally do not conduct electricity because their constituent particles are bound by strong ionic bonds in a lattice, but when in a solution, the particles are dissociated and can move easily, allowing the solution to conduct electricity
  • Solid at Room Temperature - Ionic compound assumes a solid lattice, where the cations and anions are arranged in an alternating sequence, making the structure stable
  • Hard and Brittle - Ionic compounds are generally hard because of their fixed and stable lattice, most ionic compounds are also brittle, an external force applied to the crystal may distort its lattice, make charges align, and then repel, causing the crystal to break

  
• Covalent Bond


  • Formed between nonmetals
  • Sharing of electrons between atoms and nonmetals
• Formation of a covalent bond

  1. Because nonmetal atoms have relatively similar electronegativities, they tend to attract valence electrons equally (or almost equally) and just share them to achieve an octet (or duet)

  2. Compounds that result from covalent bonding are called molecular compounds
• Type of Covalent Bond

  • Single Covalent Bond (-)
  • Double Covalent Bond (=)
  • Triple Covalent Bond
• Formal Charge


  • Covalently bonded atoms do not always share electrons equally, some atoms in a molecule or polyatomic ion have a higher electronegativity than others; thus, attract the shared electrons towards themselves greater than others
  • This results in uneven charge distribution within the molecule or ion; that is; some sites are electron-rich, electron–poor, or neutral
  • Formal charge compares the number of electrons "owned" by an atom in a molecule or ion with those possessed by the atom in a free state
  • The formula for finding the formal charge of an atom in a molecule or ion is given by: Formal Charge = Valence Electrons - Nonbonding Electrons - 1/2 Bonding Electrons
• Molecular Geometry


  • Describes the 3- dimensional arrangement of atoms within a molecule or polyatomic ion
  • The molecular geometry of molecules or ions that contain only a few atoms can be predicted using the molecule's Lewis structure and the valence shell electron pair repulsion (VSEPR) theory
  • The VSEPR theory suggests that electron pairs around an atom assume an arrangement in space that reduces the repulsions between them
  • This arrangement depends on the number and type of electron pairs (whether bonding or nonbonding)
  • The electron domain (ED) geometry is not necessarily the molecular geometry, the molecular geometry only looks at the arrangement of the atoms; the ED geometry considers the effect of the nonbonding domains on the shape of the molecule or ion
  • Since electron domains tend to repel each other, the ideal arrangement of atoms in a molecule or ion is the one that minimizes electron pair repulsion

  
• Formal charge of an atom


  Formula for finding the formal charge of an atom in a molecule or ion

  
• Molecular geometry


  • Describes the 3-dimensional arrangement of atoms within a molecule or polyatomic ion
  • Can be predicted using the molecule's Lewis structure and the valence shell electron pair repulsion (VSEPR) theory

  
• VSEPR theory


  Suggests that electron pairs around an atom assume an arrangement in space that reduces the repulsions between them
• The electron domain (ED) geometry is not necessarily the molecular geometry
• Nonbonding domains tend to spread out and occupy a larger space than a bonding domain
• Arrangement of electron domains

  • Two Electron Domains – separated by 180̊
  • Three Electron Domains – separated by 120̊
  • Four Electron Domains – separated by 109.5̊
• Polarity


  A separation of electric charges leading to a molecule or its chemical groups having an electric dipole moment, with a negatively charged end and a positively charged end

  
• Polar molecules must contain polar bonds due to a difference in electronegativity between the bonded atoms
• A polar molecule always contains one or more polar bonds; but some molecules with polar bonds can be nonpolar overall
• Kinetic Molecular Theory


  Describes the states of matter in terms of arrangement of particles, kinetic energy of particles, particle motion, attractive forces between particle and intermolecular forces

  
• Intramolecular forces


  Bonding or intramolecular forces exist inside the molecule and are relatively strong because their charges are larger and closer
• Types of intramolecular forces

  • Ionic
  • Covalent
  • Metallic

  
• Intermolecular forces


  Also known as Van der Waals force, bonding occurs between or among molecules, are relatively weak because they involve smaller charges that are farther apart
• Dipole-dipole and London dispersion forces are collectively called Van der Waals forces
• The physical properties such as melting point, boiling point, vapor pressure, evaporation, viscosity, surface tension, and solubility are related to the strength of attractive forces between molecules