Physical science

• Quantum numbers describe the characteristics of electrons and their orbitals
• An orbital is a three-dimensional region surrounding the nucleus and represents the probable location of the electrons
• Remember that every atom has a certain number of electrons, which orbit the nucleus
• Principal Quantum Number (n):

  • Indicates the size of the orbital
  • The bigger the n is, the greater is the average distance of an electron
  • Also indicates the main energy level occupied by an electron
  • Can easily be determined based on the period (row) the atom is located in the periodic table
• Azimuthal Quantum Number (l):

  • Also known as angular momentum quantum number
  • Represents the shape of the orbital
  • Allowed values: l= 0 to n-1
  • Each value represents the type of orbital: s, p, d, f orbitals
• Magnetic Quantum Number (ml):

  • Indicates the orientation of an orbital around the nucleus
  • Possible values: 2l+1; integers from –l to +l, including 0
  • For example, if l=0, only one value for ml is possible, that is ml=0
  • If l=1, there are 3 possible values of ml: -1, 0, and +1
• Spin Quantum Number (ms):

  • Indicates the spins of the electrons and may only have 2 possible values, +1/2 and -1/2
  • The signs only refer to the orientation of the spins, not on the electric charge
  • The orientation is usually upward or downward when represented in diagrams
• Electron Configuration uses the symbols of the orbitals and the number of electrons (written as superscripts) that occupy each orbital
• Orbital Diagram consists of boxes and arrows that represent the orbitals and the electrons, respectively
• Three general rules in electron distribution:

  • Aufbau principle: electrons should occupy first the orbitals with lower energy before those with higher energy
  • Pauli exclusion principle: no two electrons in an atom can possess the same set of quantum numbers
  • Hund’s rule of maximum multiplicity: each orbital in a subshell is singly occupied before pairing of electrons occurs
• Noble Gas Electron Configuration:

  • Long electron configurations can be shortened using core symbols
  • Core symbols are representations of the electron configuration of the noble gas that belongs to the row before that of the element
• Atoms are more stable when bonded with other atoms in a compound
• At the end of the lesson, students should be able to:

  • Define electronegativity
  • Recognize the electronegativity of elements in the periodic table
  • Predict the bonding between atoms based on electronegativity difference
• Lewis Dot Symbols:

  • In chemical reactions, the electrons in the outermost shell are crucial
  • Lewis dot symbols visually emphasize the outermost electrons of elements
• Valence Electrons:

  • Valence electrons are found in the outermost shell of an orbital
  • Valence electrons participate in chemical reactions
• The Octet Rule:

  • Atoms 'want' to fill their electron shells completely
  • The Octet Rule states that atoms prefer to have eight electrons in the valence shell
• Chemical Bond:

  • Is an electrical attraction between the nuclei and valence electrons of an atom
  • Binds atoms together
  • Three types of bonds: ionic, covalent, and metallic
• Ionic Bonds:

  • Formed between a metal and a nonmetal due to a large electronegativity difference
  • Result from the transfer of valence electrons from one atom to another
• Covalent Bonds:

  • Occur when atoms share electrons to fill their energy shells
  • Happen between nonmetals
  • Can be polar covalent or nonpolar covalent
• Metallic Bonds:

  • Bonds holding metal atoms together are called metallic bonds
• Electronegativity:

  • Measure of an atom's tendency to attract electrons towards itself
• The absolute value of electronegativity difference (∆EN) between two atoms determines the type of chemical bond between them:

  • Ionic bond: ∆EN > 1.7
  • Polar covalent bond: 1.7 > ∆EN > 0.4
  • Nonpolar covalent bond: ∆EN < 0.4