Trends of the Periodic Table

1. Trends of the Periodic Table

2. Review! • Periodic Table was first organized by… • Dmitri Mendeleev in the mid 1800’s • Mendeleev organized the elements by chemical reaction in rows, then by atomic mass in columns • Henry Moseley then took Mendeleev’s table, kept the chemical reactivities together, but placed them in columns instead. He also ordered the elements by increasing atomic number in rows. • When Moseley did this, all the periodic trends just fell into place. • Remember: columns = groups/families, rows = periods

3. Periodic Trends

4. Electrons • Electrons do not freely float in space • Orbit around nucleus: Electron shells • Each shell corresponds to an amount of energy.

5. Valence Electrons • The valence electrons are the outermost electrons of an atom. • The valence electrons determine the chemical properties • Number of valence electrons equals the column number in the “A” columns • Elements with the same number of valence electrons are very similar chemically – Alkali metals in Group 1A – 1 valence electron Li, Na, K, Rb, Cs – Halogens in Group 7A – 7 valence electrons • F, Cl, Br, I

6. Atomic Radius • What is Atomic Radii? • Distance from the nucleus to the outermost level of e- (aka the valence shell) • What trend do you see as you go across (left to right) the period? • Atomic radius decreases • Down the group? • Atomic Radius increases • WHY???

7. Explaining the Trend • As you go L to R, the atomic radius decreases because as you go L to R, the amount of attraction between p+ and e- increase. More attractions = smaller atomic radius • As you go down a column, atomic radius increases because the e- are farther away from the nucleus. There are weaker attractions. Weaker attractions = larger atomic radius

8. Electronegativity • What is Electro-negativity? • An atom’s Luuuvvv for electrons! • The tendency to attract another atom’s electrons • What trend do you see as you go across the period? • Electronegativity increases! • Down the group? • Electronegativity decreases! • WHY???

9. Explaining the Trend • As you go L to R, electronegativity increases because of the increase in protons. The more protons, the more able it will be to attract other atom’s electrons. More attractions (small radius) = large electronegativity • As you move down a column, electronegativity decreases because of the increase in number electron an atoms already has. This means the atom will be less able to attract another atom’s electrons. • Less attractions (large radius) = small electronegativity

10. Ionization Energy • What is Ionization Energy? • The energy needed to remove an electron • What trend do you see as you go across the period? • Ionization E increases • Down the Group? • Ionization E decreases • WHY???

11. Explaining the Trend • As you go L to R, the ionization energy increases because of the increase in the number of protons. The more protons, the more energy that is needed to remove an electron. More attractions (small radius) = large ionization energy • As you go down a column, the ionization energy decreases because of the decrease in attractions. • Due to electron shielding • More electrons, leads to outer electrons less tightly held. • The less attractions, the lower the energy that is needed to remove an electron. Less attractions (large radius) = small ionization energy

12. Ionization Energy • Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. • First ionization energy is that energy required to remove first electron. • Second ionization energy is that energy required to remove second electron, etc.

13. Ionization Energy • It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionization energy takes a quantum leap.

14. Electron Affinity • What is Electron Affinity? • The energy needed to add an electron • As you go across the period electron affinity increases . • Electron affinity decreases down the family • WHY???

15. Explaining the trend • As you go L to R, the electron affinity increases because of the increase in the number of protons. The more protons, the greater the attraction the protons have for electrons. More attractions (small radius) = large electron affinity • As you go down a family, the electron affinity decreases because of the decrease in attractions. • Due to electron shielding • More electrons, leads to outer electrons less tightly held. • The less attractions, the lower the electron affinity Less attractions (large radius) = small electron affinity

16. Homework • Worksheet(s)