The Periodic Table

1. The Periodic Table

2. Periodic Table Groups to Know…

3. The Periodic Table • The elements are arranged in rows and columns on the Periodic Table according to the similarities in their properties • Metals are found on the left side and bottom of the Periodic Table • Non-metals are found on the right side of the Periodic Table

4. The Periodic Table • Metalloids are arranged in a diagonal pattern between the metals and non-metals

5. The Periodic Table • A column of elements is known as a GROUP • The group number corresponds to the number of valence electrons • A horizontal row of elements is known as a PERIOD • The period number corresponds to the number of shells containing electrons

6. The Periodic Table • Four groups have characteristic family names : • Group I Alkali metals • Group II Alkaline earth metals • Group VII Halogens • Group VIII Noble (inert) gases

7. The Periodic Table • The large block of elements between groups 2 and 3 forms a transition between the metals on the left and non-metals on the right and are called the Transition metals • Elements 58-71 are called Lanthanides (or rare earth elements) because of their rare occurrence in nature

8. The Periodic Table • Elements 90-103 are called Actinides and these are all radioactive elements • Elements with atomic numbers more than 92 are called transuranium elements. They are not found in nature and are synthesised in the laboratory

9. Group I: Alkali metals • Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Caesium (Cs) and Francium (Fr) • Each has one electron in the valence shell • Lose the one electron in chemical reactions to form the M+ ion

10. Group I: Alkali metals • Very reactive with water (stored under paraffin oil to exclude moisture) • Form white crystalline solids which are soluble in water

11. Group II: Alkaline Earth metals • Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), Radium (Ra) • Each has two electrons in its valence shell • Lose 2 electrons to form M2+ ions in chemical reactions

12. Group II: Alkaline Earth metals • Less reactive with water than group I metals • From white crystalline solids, some of which are soluble in water

13. Group IV • Carbon is a non metal, Silicon and Germanium are metalloids and Tin and Lead are metals • 4 valence electrons • Carbon, silicon and germanium react to form compounds of general formula XY4 where X is the group IV element • Tin and lead form +2 and +4 ions

14. Group VII: Halogens • Exist as diatomic molecules: Fluorine F2 Chlorine Cl2 Bromine Br2 Iodine I2 Astatine At2 • Are very reactive non metals • Occur in nature as the halide ions X- in salts when reacted with metals (i.e. they have gained one electron)

15. Group VIII: The Noble Gases • Helium He Argon Ar Neon Ne Krypton Kr Xenon Xe Radon Rn • Are unreactive non metals due to their stable electronic configuration • Do not usually form compounds (He and Ne do not form compounds)

16. The Transition Elements • Located between Group II and III • Have relatively high melting points and densities • Usually form coloured ions in solutions Copper: blue Iron(II): green Iron(III): brown Chromium: orange Nickel: green

17. Trends in the Periodic Table • Elements within a group have the same number and distribution of valence electrons • = some patterns/trends in behaviour • Reactivity • Atomic Radius • Electronegativity

18. Trend: Reactivity • Most reactive: • Alkali Metals  lose one electron to from an octet • Halogens  gain one electron to from an octet

19. Trend: Reactivity Increasing Increasing Increasing Increasing

20. Trend: Atomic Radius • Atomic Radius = a measurement of the size of an atom (radii = plural) • Atomic radii increases as we move down a group • More electron shells = bigger • With each additional energy level the attractive forces of the nucleus (protons, +) is reduced = shielding effect • Valence electrons (outer shell) are not as strongly attracted = bigger radius

21. Trend: Atomic Radius • Atomic radii decreases from left to right across a period • Moving from left to right atoms have more protons • Attractive forces of the nucleus (+) for valence electrons (-) increase • Electrons pulled closer to the nucleus = smaller radius

22. Increasing Increasing

23. Trend: Electronegativity • Electronegativity= a number that describes the ability of an atom to attract electrons • Provides a means of prediction the nature of the forces that hold a pair of atoms together in a bond • More electronegative = stronger pull on electrons being shared • Less electronegative = weaker pull on electrons being shared

24. Trend: Electronegativity Increasing Increasing

26. Atomic Radius & Electronegativity • Electronegativityincreases when the atomic radius is small • Electronegativitydecreases when the atomic radius is large • Large atomic radius = valence electrons are further from the nucleus • Less attractive forces holding the valence electrons • Less attractive forces to pull electrons to t he atom