The Periodic Table

1. The Periodic Table

2. Development of Periodic Table Dmitri Mendeleev grouped elements by similar properties.

3. Mendeleev’s Table 1869

4. Development of Periodic Table Mendeleev predicted the discovery of germanium (which he called eka-silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.

5. Metals, Non-metals and Metalloids

6. Development of Periodic Table • Elements are arranged in vertical columns called GROUPS and horizontal rows called PERIODS.

7. Periodic Table • Elements in the same group have similar properties because they have the same number of electrons in their outermost shell. • These electrons are called VALENCE ELECTRONS.

8. Significance of Groups • Group 1 elements all have one electron in their outer shell • Group 2 elements all have two electrons in their outer shell • These electrons give elements in the same group similar properties.

9. Significance of Periods • The elements in the first period all have one shell. • The elements in period 2 all have 2 shells. • The elements in period 3 all have 3 shells • And so on

10. Hydrogen Group • This element does not match the properties of any other group, so it stands alone. It is place above group 1, but it is not part of that group. • It is very a reactive, colorless, odorless gas at room temperature.

11. Group 1 – Alkali Metals • Alkali metals are extremely reactive and are never found in nature in their pure form. They are silver colored and shiny. Their density is extremely low so that they are soft enough to be cut with a knife. (1 outer level electron)

12. Group 2 – Alkaline Earth Metals • Alkaline-earth Metals – Slightly less reactive than alkali metals. They are silver colored and harder and denser than alkali metals. (2 outer level electrons) • Less reactive than group 1, but are also too reactive to be found in nature as free elements.

13. Groups 3 – 12 – Transition Metals • Transition Metals have a moderate range of reactivity and a wide range of properties. In general, they are shiny and good conductors of heat and electricity. They also have higher densities and melting points than groups 1 & 2. (1 or 2 outer level electrons)

14. Lanthanide and Actinides • The elements in each of these two periods share many properties. The lanthanides are shiny and reactive. The actinides are all radioactive and are therefore unstable. Elements above 92 do not exist in nature but have been manufactured in the lab.

15. Lanthanide and Actinides • Periodic table with lanthanide and actinides in their “natural” position.

16. Group 17 - Halogens • Halogens are all nonmetals. Very reactive and poor conductors of heat and electricity.

17. Group 18 – Noble Gases • Noble Gases are unreactive nonmetals. All are colorless, odorless gases at room temperature. All found in earth’s atmosphere in small amounts. (8 outer level electrons) • The outer electron shell is full making them unreactive.

18. Groups 13 - 16 • Group 13: Boron Group –Reactive. (3 outer level electrons) • Group 14: Carbon Group –Varied reactivity. (4 outer level electrons) • Group 15: Nitrogen Group –metal. Varied reactivity. (5 outer level electrons) • Group 16: Oxygen Group – Reactive group. (6 outer level electrons)

19. Valence Electrons • Valence Electrons are the electrons in the outermost shell. • Group 1 (1A) – 1 valence electron • Group 2 (2A) – 2 valence electrons • Group 13 (3A) – 3 valence electrons • Group 14 (4A) – 4 valence electrons • Group 15 (5A) – 5 valence electrons • Group 16 (6A) – 6 valence electrons • Group 17 (7A) – 7 valence electrons • Group 18 (8A) – 8 valence electrons.

20. Valence Electrons • For groups 3 – 12 (transition metals) as well as the lanthanide and actinide series the number of valence electrons can not be easily determined from the periodic table.

21. Sizes of Atoms The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

22. Sizes of Atoms Bonding atomic radius tends to… …decrease from left to right across a row. …increase from top to bottom of a column due to increasing value of n

23. Ions • An ion is an atom or group of bonded atoms that has a positive or negative charge • A positive ion is a cation. • A negative ion is an anion.

24. Sizes of Ions • Ionic size depends upon: • Nuclear charge. • Number of electrons. • Orbitals in which electrons reside.

25. Sizes of Ions • Cations are smaller than their parent atoms. • The outermost electron is removed and repulsions are reduced.

26. Sizes of Ions • Anions are larger than their parent atoms. • Electrons are added and repulsions are increased.

27. Sizes of Ions • Ions increase in size as you go down a column. • Due to increasing value of n.

28. Ionization Energy • Amount of energy required to remove an electron from the ground state of an 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.

29. Trends in First Ionization Energies • As one goes down a column, less energy is required to remove the first electron.

30. Electronegativity: • The ability of atoms in a molecule to attract electrons to itself. • On the periodic chart, electronegativity increases as you go… • …from left to right across a row. • …from the bottom to the top of a column.