Explaining Periodic Trends Textbook Pages: 31-40

Periodic Trends O Atomic Radius O Ionization Energy O Electron Affinity O Electronegativity O When examining Periodic Trends ALWAYS look at: O Number energy levels O Number of protons O As we go down a group: O Energy levels increase O As we go across a period (L to R): O Number of protons increases

Atomic Radius O Definition: The distance from the center of an atom to the boundary within which the electrons spend 90% of their time. O Ex: Which atom would be larger: Be or Mg? Explain. O Be has 2 energy levels O Mg has 3 energy levels O Therefore, Mg is larger O Ex: Which atom would have the smallest radius: Mg or Si? Explain. O Mg and Si both have 3 energy levels O Si has more protons to attract the electrons O Therefore, Si is smaller (smallest radius)

Atomic Radius O In General: O Down a Group: O Atomic Radius increases (more energy levels) O Across a Period (L to R): O Atomic Radius decreases (same energy levels, more protons) O The atomic radius of Bromine is larger than that of Nitrogen. Why do you think this is so?

Ionization Energy O Definition: The amount of energy required to remove an electron from the outermost energy level of an atom or ion (in the gaseous state). O Ex: Which atom has the larger ionization energy: F or O? Why? O Both F and O have the same energy levels O F has more protons to attract the same number of energy levels O F will hold the electrons more tightly O Therefore, F has a larger ionization energy

Ionization Energy O Ex. Would more energy be required to remove an electron from a Ne atom or from a F ion? O Both F and Ne have the same number of energy levels O Ne has more protons to attract the electrons, making them more difficult to remove O Therefore, Ne would require more energy to remove an electron (higher ionization energy) O Ex: Which atom has the smallest ionization energy: Li or Rb? Explain. O Li has 2 energy levels O Holds its electrons more tightly because they are closer to the nucleus O Rb has 5 energy levels O Takes less energy to remove the outermost electron because it is farther from the nucleus O Therefore, Rb has the smallest ionization energy

Ionization Energy O In General: O Down a Group: O Ionization Energy decreases (more energy levels) O Across a Period (L to R): O Ionization Energy increases (same energy levels, more protons)

Electron Affinity O Definition: Ability of an atom to attract electrons. O In General: O Down a Group: O Electron Affinity decreases (more energy levels) O Across a Period (L to R): O Electron Affinity increases (same energy levels, more protons)

Electronegativity O Definition: Ability of an atom to attract electrons in a bond. O In General: O Down a Group: O Electronegativity decreases (more energy levels) O Across a Period (L to R): O Electronegativity increases (same energy levels, more protons)

Ionization Equations O Atoms will lose or gain electrons so that they are isoelectronic (have the same number of electrons) with the nearest noble gas O This makes the atoms stable because their orbits are full (stable octet) O Examples: O F + e−  F− O Mg  2e− + Mg 2+ O Isoelectronic with Ne (noble gas)

Summarizing Trends in the Periodic Table O Take a look at Page 38 in your textbook!!!