1. SL Topic 3 Periodic Trends Wichita East High School Beth McKee The Periodic Table – p. 11 IB Diploma book

2.  Arranged by atomic number so that there are repeating trends.  Groups – Elements in the same column that have the same number of valence electrons. Also called _____________.  Group 1 – alkali metals  Group 2 – alkaline earth metals  Group 7 – halogens  Group 8 or 0 – noble gases, also called the rare gases or inert gases  Periods – horizontal rows that have the same number of energy levels, the same outer shell of valence electrons.

3. Periodicity  Elements in the same group tend to have similar chemical and physical properties.  Across a period, there is a change in physical and chemical properties.  Repeating pattern in atomic radii, ionic radii, ionization energies, electronegativities and melting points.

4. Atomic Radius Atomic radius – half the distance between two nuclei of the same element. Why is it measured this way?

5. Trends in Atomic Size Group Trends – atoms get bigger as they go down a family because they are adding energy levels Periodic Trend – atoms get smaller as you go across the period from left to right. The reason is that you are adding the electrons in the same energy level and the nuclear charge is increasing with additional protons. This pulls the valence electrons in closer to the atom.

6. Trends in Atomic Size

7. Trends in Atomic Size

8. Trends in ionic size Group trends – they get larger as you go down a family, more energy levels Periodic trends – –Cations – smaller than the atoms because they lose electrons and they get smaller as you lose more electrons. It also has lost its outermost energy level, which decreases the size.

9. Trends in ionic size Anions – larger than the atom they form from because they gain electrons and they get smaller as they gain fewer electrons. Isoelectronic – ions of different elements that contain the same number of electrons so that they are like the noble gases. As protons increase, but electrons remain constant, size goes down.

10. Trends in ionic size

11. Start Ch. 6 ? #57, 59,61, 63, 76, 77, 78

12. Trends in Melting Points The changing of a solid to a liquid, particularly of metals Dependent on both the structure of the element and on the type of attractive forces holding the atoms together. Periodic trend – increases across the metallic period, with metallic bonding, as there are more valence electrons to hold the metal together. You have more electrons to move, therefore it takes more energy.

13. Con’t periodic trends – silicon and carbon form macromolecular covalent structures with very strong bonds and thus very high melting points. The nonmetals have simple molecular structures with weak van der Waals’ forces have lower melting points (some are gases at room temp.) Group trend – Group 1 decreases down a family as they get larger and metallic bonding decreases. Group 7 – attractive forces increase as the diatomic molecules get larger, so melting points go up.

14. Trends in ionization energy The energy required to remove an electron from an element in its gaseous state »X(g)  X + (g) + e - Group trend – decreases as you go down a family – they are getting ___________ Periodic trend – it increases as you move across the period because the atoms are getting smaller and the valence electrons are closer/held more tightly by the nucleus.

15. Electronegativity The relative measure of attraction that an atom has for a shared pair of electrons when it is covalently bonded to another atom. As the size of the atom decreases the electronegativity increases, so the value increases across a period and decreases down a group. Nitrogen, oxygen and fluorine have the highest electronegativities which plays a large part in hydrogen bonding between molecules. It is a composite number developed by Linus Pauling from several measurements such as size and ionization energy.

16. Electron Affinity The energy change associated with gaining of an electron.

17. The Periodic Table and Chemical Properties Group 1 – the alkali metals They have increasing reactivity down the family since they have one valence electron that they readily lose. As successively higher energy levels are added, less energy is required to remove the valence electron. This makes them very good reducing agents.

18. They are called alkali metals because they all react with water to form an alkali solution of the metal hydroxide and hydrogen gas. Lithium floats and reacts quietly. 2Li(s) + 2H 2 O(l)  2Li + (aq)+ 2OH - (aq) + H 2 (g) Sodium melts into a ball which darts around on the surface. 2Na(s) + 2H 2 O(l)  2Na + (aq)+ 2OH - (aq) + H 2 (g) Potassium reacts violently as it generates heat that quickly ignites the hydrogen gas. 2K(s) + 2H 2 O(l)  2K + (aq)+ 2OH - (aq) + H 2 (g) Start this link first!

19. They also react readily with chlorine and bromine to form ionic salts 2Na(s) + Cl 2 (g)  2Na + Cl - (s) 2K(s) + Br 2 (l)  2K + Br - (s)

20. The Periodic Table and Chemical Properties Group 7 – the halogens They have decreasing reactivity down the family, as they have seven valence electrons and react by gaining one more electron to form the halide ion. As the halogen gets larger down the family the attraction for another electron decreases since the valence shell is further from the nucleus, so reactivity goes down. They are good oxidizing agents.

21. In single replacement reactions, the higher activity halogen will substitute for a lower activity halogen that is in a compound. Cl 2 (aq) + 2Br - (aq)  2Cl - (aq) + Br 2 (aq) Cl 2 (aq) + 2I - (aq)  2Cl - (aq) + I 2 (aq) Br 2 (aq) + 2I - (aq)  2Br - (aq) + I 2 (aq)

22. Test for halide ions Halide ions in solution can be detected by adding silver nitrate solution. The silver reacts with the halide ion to form a precipitate of the silver halide. The silver halides can be distinguished by their colour. These silver halides react with light to form silver metal, which is the basis for film photography. Ag + (aq) + X - (aq)  AgX(s) AgX(s) Ag(s) + ½ X 2 If X = Cl – AgCl is white Br – AgBr is cream I - AgI is yellow

23. Changing metallic to nonmetallic properties of the elements across Period 3 Metals tend to be shiny and good conductors of heat and electricity. In period 3 they are sodium, magnesium, and aluminum. Metal oxides tend to form bases in solution. Aluminum is an exception in that its’ oxide is amphoteric and can act as an acid or a base, depending on what it is reaction with. Na 2 O(s) + H 2 O(l)  2Na + (aq) + 2OH - (aq) (notice the difference between pure sodium’s reaction with water) MgO(s) + H 2 O(l)  Mg(OH) 2 (aq)

24. Changing metallic to nonmetallic properties of the elements across Period 3 Silicon is a metalloid or semimetal that is a semi-conductor as well as being shiny like metals, but it is brittle like nonmetals. Nonmetals in period 3 are phosphorus, sulfur, chlorine and argon. They do not conduct heat or electricity. Nonmetallic oxides in solution tend to be acidic. SO 3 (g) + H 2 O(l)  H 2 SO 4 (aq) sulfuric acid Cl 2 (aq) + H 2 O(l)  HCl(aq) + HClO(aq)