Periodic Trends

• Electronegativity is a measure of the ability of an atom to attract a bonding pair of electrons towards itself in a covalent bond.
• As we move down a group, electron shielding increases, leading to an increase in atomic size.
• Electron shielding is the effect that inner shell electrons have on outer shell electrons, reducing their attraction to the nucleus.
• As you move across a period, elements gain electrons and experience an increase in electron shielding, causing ionization energy to decrease.
• The atomic radius decreases across the period due to increased nuclear charge, which attracts electrons more strongly.
• Ionization energy generally decreases down a group on the periodic table due to the increasing atomic size and shielding effect.
• Atomic radius decreases down a group due to increased nuclear charge pulling in electron shells closer to the nucleus.
• The first ionization energy decreases as we go down a group due to increased screening by inner shells.
• Metals lose electrons to form positive ions (cations), while non-metals gain electrons to form negative ions (anions).
• Ionization energies are measured by removing one electron at a time until all valence electrons are removed.
• First IE is the energy needed to remove the most loosely held electron from a neutral atom.
• Electron affinity refers to the amount of energy released when an extra electron is added to a gaseous atom or ion.
• Electronegativity increases across a period, with fluorine being the most electronegative element.
• Atomic radii decrease across a period and increase down a group.
• Electron Affinity is the amount of energy absorbed (+) or released (-) when an atom gains an electron (forming an anion).
• Non-metals form stable anions after gaining an electron, and release energy (-).
• When metals gain an electron, they form unstable anions and absorb energy (+).
• The Halogens form unstable anions and absorb energy (+) because they already have a full octet and do not want to gain another electron.
• The reason why metals form unstable anions and absorb energy (+) is because they do not want to gain any electrons, instead wanting to remove them, so it requires energy to be absorbed (+) so that they can hold onto the electron. They do not have a strong attraction to the electron.
• Atomic radius is the size of an atom.
• Ionization energy is the amount of energy required for an atom to remove an electron.
• Nuclear charge is the amount of protons in an atom's nucleus, and its ability to attract electrons.
• The shielding effect occurs when electrons in between the valence shell and nucleus effectively block the attractive forces between the nucleus and the outer electrons.
• The effective nuclear charge (ENC) is the apparent charge of the electrons in orbitals as a result of the shielding effect.
• Metals reactivity decreases across a period because:

  1. Atomic radius decreases
  2. Shielding decreases
  3. Stronger attraction between the nucleus and the valence electrons; therefore the valence electrons have a harder time removing electrons
  4. Ionization energy increases b/c electrons are harder to lose
• Metals reactivity increase down a group b/c:

  1. Shielding increases
  2. Number of shells increase
  3. Weaker attraction between the nucleus and the valence electrons, therefore valence electrons are easier to remove
  4. Ionization energy decreases due to removal of electrons being easier
• Non-metals reactivity increases across a period b/c:

  1. Shielding decreases
  2. Stronger attraction between the nucleus and valence electrons, which pulls the valence electrons in, making the atom smaller
  3. Atomic radius decreases
  4. Therefore easier to gain electrons
• Non-metals reactivity decreases down a group b/c:

  1. More shells
  2. Larger Atomic Radius
  3. More shielding
  4. Weaker attraction between the nucleus and the valence electrons, therefore v. electrons are not pulled in as much, making it harder to gain electrons
• Electron affinity increases across a group:

  1. Decrease in shielding
  2. Stronger attraction b/w nucleus and valence electrons
  3. Closer to a full octet, therefore more energy is released
  4. More stable ion, nucleus has a strong hold on the valence electrons
• Electron affinity decreases down a group:

  1. More shells; More shielding
  2. Weaker attraction between nucleus and v. electrons
  3. Nucleus has a weaker hold on the new electron added
  4. Less energy is released
• Atomic radius decreases across a period:

  1. Less shielding
  2. Larger effective nuclear charge
  3. Stronger attraction between the nucleus and the valence electrons
  4. nucleus pulls in electrons making it smaller
• Atomic Radius increases down a group:

  1. More shells; more shielding
  2. Effective nuclear charge decreases
  3. Weaker attraction between the nucleus and valence electrons, therefore a weaker pull, making the atom bigger
• Ionization energy increase across a period:

  1. Less shielding
  2. Stronger attraction b/w nucleus and valence electrons, making it harder to lose an electron
  3. Atomic radius decreases
  4. More energy required to lose an electron
• Ionization energy decreases down a group:

  1. More shells; more shielding
  2. Weaker attraction b/w the nucleus and valence electrons, making it easier to lose an electron
  3. Atomic radius increases
  4. Less energy required to lose an electron
• Electronegativity increase across a period:

  1. Less shielding
  2. Stronger attraction between the nucleus and valence electrons, making it easier for the nucleus to attract electrons in a bonding pair
  3. Atomic radius decreases
• Electronegativity decreases down a group:

  1. More shells; more shielding
  2. Weaker attraction between the nucleus and valence electrons
  3. Atom does not pull/attract shared electrons as close
  4. Atomic radius increases
• Greater negative values = Greater electron affinity = More stability
• Greater attraction = greater stability