Electronic Structure and Periodicity

• Periodic table


  Tabular arrangement of the elements in order of increasing atomic number such that elements having similar chemical properties are positioned in vertical columns
• Periodic table


  • Contains the symbol, atomic number and atomic mass of each element
  • No information concerning neutrons is available
  • Location of element is specified by period number and group number
• Period


  Horizontal row of elements in the periodic table
• Group


  Vertical column of elements in the periodic table
• Notations for periodic table groups

  • Roman numerals and the letters A and B
  • Arabic numbers (1 through 18)
• Classification of elements


  • Based on selected physical properties (metals, nonmetals, metalloids)
  • Based on electron configurations (noble-gas, representative, transition, inner transition)
• Metals


  Elements with the characteristic properties of luster, thermal conductivity, electrical conductivity, and malleability
• Nonmetals


  Elements characterized by the absence of the properties of luster, thermal conductivity, electrical conductivity, and malleability
• Metalloids


  Elements which have the properties of both metals and nonmetals
• Selected physical properties of metals and nonmetals

  • Metals: High electrical conductivity, high thermal conductivity, metallic luster, mostly solids, malleable, ductile
  • Nonmetals: Poor electrical conductivity, good heat insulators, no metallic luster, solids/liquids/gases, brittle, nonductile
• Periodic law


  When "elements are arranged in order of increasing atomic number", elements with similar chemical properties occur at periodic (regularly recurring) intervals
• Valence electrons


  Electrons found in the outermost shell that interact when atoms form bonds
• Valence electrons for different element types

  • Representative and noble-gas elements: In s and p subshells of highest energy level

  Transition elements: In outermost energy level and d orbitals of next inner main energy level

  Inner transition elements: Outermost s, d and f electrons
• Effective nuclear charge (Zeff)


  The "nuclear charge felt by an electron" when both the actual nuclear charge (Z) and the repulsive effects (shielding) of the other electrons are taken into account
• Across a period


  Effective nuclear charge increases
• Down a group


  Effective nuclear charge also increases, but the increase is smaller than across a period
• Effective nuclear charge (Zeff)


  The nuclear charge experienced by an electron when both the actual nuclear charge (Z) and the repulsive effects (shielding) of the other electrons are taken into account
• Trends in Zeff for valence electrons in the periodic table

  1. From left to right across a period, the effective nuclear charge increases

  2. From top to bottom or going down a column, effective nuclear charge also increases
• Atomic radius


  One-half the distance between the two nuclei in two adjacent metal atoms or in a diatomic molecule
• Trends in atomic radius


  1. left to right across a period, atomic radius decreases

  2. top to bottom of a group, the atomic radius increases
• Ionization energy


  The minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion
• Trends in ionization energy


  1. left to right across a period, ionization energy generally increases with increasing atomic number

  2. top to bottom of a group, ionization energy generally decreases with increasing atomic number
• The trend in ionization energy is not followed when the added valence electron in the next element enters a new sublevel (higher energy sublevel) or is the first electron to pair in one orbital of the sublevel (electron repulsions lower energy)
• Electron affinity


  The energy change that occurs when an "electron is added to a gaseous" atom to form negatively charged ions
• Trends in electron affinity


  1. There is a general correlation between electron affinity and effective nuclear charge, which also increases form left to right in a given period

  2. Electron affinities do not change greatly as we move down a group
• The electron affinity of noble gas is positive since the addition of an electron to a noble gas requires that the electron reside in a higher-energy subshell that is empty and is energetically very unfavorable
• Halogens (Group 7A) have the most negative electron affinities since by gaining an electron, they will form a stable negative ion that has a noble-gas configuration
• Electron affinities of the Group 5A elements are either positive like Nitrogen or less negative than their neighbors to the left because these elements have half-filled p subshells, and the added electron must be put in an orbital that is already occupied, resulting in larger electron-electron repulsion
• As we proceed from F to I in the halogens, the average distance between the added electron and the nucleus steadily increases, causing the electron-nucleus attraction to decrease, but the orbital that holds the outermost electron is increasingly spread out, so the electron-electron repulsions are also reduced, resulting in the reduction in the electron-nucleus attraction being counterbalanced by the reduction in electron-electron repulsions