PERIODIC TRENDS

• Dmitri Mendeleev, a Russian chemist, perfected the periodic table. Element 101 is named after him
• Mendeleev predicted elements that had not yet been discovered based on their similarities and patterns
• Elements with similar chemical properties and reactions have a similar valence electron in their highest energy level
• Chemical families in the periodic table are known as groups
• Groups in the periodic table:

  • Group 1: Alkali Metals
  • Group 2: Alkaline Earth Metals
  • Group 3-12: Transition Metals
  • Group 17: Halogens
  • Group 18: Noble Gases
• Lanthanides and Actinides are part of the periodic table
• Information on valence electrons and properties of an element are repeated in certain trends
• Regular increments in certain trends led to the development of the Periodic Law
• Valence electrons increase from left to right in the periodic table
• Atomic number increases from left to right in the periodic table
• Within a group, elements get larger as you go down
• Size of atoms increases as you go down a group
• Other trends in the periodic table include:

  • Effective nuclear charge
  • Atomic radius
  • Ionic radius
  • Electron affinity
  • Electronegativity
  • Ionization energy
• Effective nuclear charge is lower due to the distance from the nucleus and shielding
• Atomic radius is the distance from the nucleus to the outermost electron, measured in picometers (2×10^-12 meters)
• Valence electrons are further from the nucleus due to shielding by inner electrons, which blocks the full attraction of the nucleus
• Repulsion from other electrons leads to less effective nuclear charge
• Metals with smaller ionic radii, such as lithium or sodium, readily lose electrons to form cations, whereas metals with larger ionic radii, such as cesium or rubidium, are more stable and less reactive.
• The ionic radius increases as you move across a period in the periodic table from left to right due to increasing electron shells.
• The atomic radius decreases across the period due to increasing nuclear charge.
• The atomic radius decreases across the period.
• Electron affinity is the energy change when an electron is added to a gaseous atom or ion, forming a negative ion.
• The general trend of decreasing ionic radius as you move down a group in the periodic table is due to increasing nuclear charge as you increase in atomic number.
• Atomic radii decrease as we move from left to right along a period due to an increase in nuclear charge, which attracts electrons more strongly towards the nucleus.
• Electron shielding is the effect that inner shells have on outer shell electrons, reducing their attraction to the nucleus.
• The atomic radius decreases across periods because there are fewer shells and subshells available for electrons to occupy, leading to increased repulsions between valence electrons.
• Metals have positive electron affinities as they gain electrons easily.
• Electron affinity refers to the energy change when an atom gains one electron to become a negatively charged ion.
• In general, elements on the left side of the periodic table have higher positive charges than those on the right side.
• As we move down groups, the number of protons increases but the number of electrons also increases, resulting in an increase in size.
• Electronegativity is the ability of an atom to attract shared pairs of electrons towards itself when it forms a chemical bond.
• Atomic size is related to the number of protons (nuclear charge) and the number of electrons.
• Increasing atomic mass causes increasing atomic radius.
• Electronegativity is the ability of an element to attract shared pairs of electrons towards itself in a chemical bond.
• As we go down a group, there is an increase in atomic size because the outermost energy level expands to accommodate additional electrons.
• Electronegativity increases down a group (column) on the periodic table.