1. Chapter 3 Notes: Periodic Trends
Chapter 3.2: Elements show trends in their physical and chemical properties across periods and down groups.
Chapter 3 Notes: Periodic Trends

2. Important terms for this section
Periodicity
Atomic radii
Ionic radii
Electronegativity
Electron affinity
Ionization energy
Melting points
Effective nuclear charge
Chemical properties
Noble gases
Alkali metals
Halogens
Period 3 oxides

3. Effective Nuclear Charge
The attraction “experienced” by the outer electrons from the positively charged nucleus
Shielding occurs moving down groups from inner electrons ~ same ENC
More p+ + e- means higher ENC moving across period

4. Effective Nuclear Charge

5. ENC continued: Transition metals
Nuclear charge increases from Scandium to copper (e- added to inner 3d sub-level)
For the most part, 3d electrons shield 4s electrons from increasing nuclear charge
The 4s electrons of the 1st row of transition metals feels only slightly increasing ENC as atomic number increases
This means the atomic radii only gradually increases as well
Increase in nuclear charge is ~equal to the inc in shielding effect  ENC remains about constant

6. ENC of first 3 rows

7. Atomic radius – Values in Table 8 of booklet
Measured: half distance between neighboring nuclei
Since ENC increases across period, radius decreases
Since more e- repulsion moving down (add an E level), radius increases

8. Ionic radius – 5 important points
Cations are smaller than parent atoms
Anions are larger than parent atoms
Ionic radii decrease from groups 1-14 for cations
Increase ENC, results in decrease of size
Ionic radii decrease from groups for anions (for same reason as above)
Positive ions are smaller than negative ions of similar electron config due to the loss of E level
Ionic radii increase down a group b/c E level increases

9. Ionic radii trends

10. Ionization energies
Energy required to remove 1 mol of e- from 1 mol of gaseous atoms in their ground state
Measure of attraction between nucleus and its outer electrons
IE generally increases across period
Increase of ENC causes inc of attraction and so difficult to remove
IE decreases down group
E- removed is furthest from nucleus
ENC inc, but more repulsion of inner e- with outer e- (shielding)
Differences in trend shows evidence of sub-levels

11. Ionization energy

12. Electron affinity – Values in Table 8 of booklet
Energy change when 1 mol of e- is added to 1 mol of gaseous atoms to form 1 mol of gaseous ions
As e- is added to pos. nucleus, usu. exothermic
Releases energy b/c the attraction is favorable
Additional e- added is endothermic due to repulsive nature of the already added e-
Thought of as the “negative” of 1st IE values
Closely related to electronegativity but not same
ENC increases, electron affinity increases

13. Electron affinity

14. Electronegativity – Values in Section 8 of booklet
The ability of an atom to attract electrons in a covalent bond
So attraction of + nucleus with bonding e- on another atom(s)
General same trend as with IE
Increases across period
Inc of ENC, so increase attraction to e-
Decreases down group
Greater distance, so decrease attraction to e-

15. Electronegativity

16. Metals vs Non-metals
Metals have lower ionization energies and electronegativities than non-metals
The ability of metals to have mobile electrons (which is why they can conduct) also demonstrates this property
Electrons that do this are called delocalized
“Sea of electrons” is the way metals bond (ch 4)

17. Melting Points
Moving down group 1, MP decreases b/c the attraction of nucleii with delocalized e- decreases with distance
Moving down group 17, intermolecular forces (IMFs), specifically London dispersion forces, increases, so MP increases
Across periods
MP increases from group 1-14
Group 15-18, drop and decrease to minimum

20. Chemical Properties: Noble Gases – Group 18
Colourless gases
Monatomic
Very unreactive
All have stable octets
He has stable first principle E level
Highest IE and almost no electronegativity
They are “full” of themselves and do not react with others due to their nobility
Note: chemical properties of elements are mainly determined by the number of valence electrons (thus, groups have similar properties)

21. Chemical Properties: Alkali Metals – Group 1
Silvery color and too reactive to be found in nature (in pure form)
Form singly charged cations = M+
Low ionization energies
Reactivity increases down group
Ability to conduct based upon outer e-

22. Chemical Properties: Alkali Metals – Reaction with H2O
Alkali metal + H2O  H2 + metal hydroxide
Increasing intensity of reaction moving down the group: Li < Na < K < Rb < Cs
Example reaction:
2K(s) + 2H2O(l)  2KOH(aq) + H2(g)
The KOH is ionic and will dissociate in water:
2K(s) + 2H2O(l)  2K+(aq) + 2OH-(aq) + H2(g)

23. Chemical Properties: Halogens – Group 17
Diatomic molecules = X2
Accepts an electron very easily
High ENC (~ +7) and so pull on e- in other atoms
Reactivity decreases down group
F2 and Cl2 are gases; Br2 is a liquid; I2 is a solid at room temp

24. Chemical Properties: Halogens – Reaction with Group 1
Very reactive with alkali metals to form ionic halides (salts)
Complementary reactants: Na+ and Cl-
2Na(s) + Cl2(g)  2NaCl(s)
High ENC of Cl2 pulls outer Na e- and electrostatic attraction completes transfer

25. Chemical Properties: Halogens – Displacement RXNs
Which is most reactive? Br or F? Cl or I?
What is the order of increasing reactivity?
What do you think would happen if you pour brominated water into chlorinated water?
What do you think would happen if you pour chlorinated water into brominated water?

26. Chemical Properties: Halogens – Displacement RXNs
When a more reactive halogen is added to a solution of less reactive:
The less reactive is pushed out of solution
Chlorine is green/yellow
Bromine is orange/brown
Iodine is violet
Halogens are more soluble in non-polar solvents

27. Chemical Properties: Halogens – Identifying Halides
When reacting with silver, halides form insoluble salts
Adding a halide solution with silver produces a precipitate
Ag+(aq) + X-(aq)  AgX(s)

28. Chemical Properties: Period 3 Oxides – Bonding
Na  Na2O
Mg  MgO
Al  Al2O3
Si  SiO2
P  P4O10
 P4O6
S  SO3
 SO2
Cl  Cl2O7
 ClO
Moving across period:
Ionic bonding transitions to covalent:
NaAl have giant ionic struc.
PCl have molecular covalent
Si has giant covalent (network)
Oxides only conduct in liquid form (moving of ions)
Oxidation number related to group number

29. Chemical Properties: Period 3 Oxides – Bonding

30. Chemical Properties: Period 3 Oxides – Acid-Base Character
Acid-base character is based upon the bonding and structure
Acid- has pH of 6 or less (covalent)
Base- has pH of 8 or more (ionic)
Amphoteric- can react with acids and bases

31. Chemical Properties: Period 3 Oxides –Basic Oxides
Basic oxides – Na and Mg dissolve in H2O to form alkaline solutions (has hydroxide ions)
Na2O(s) + H2O(l)  2NaOH(aq)
MgO(s) + H2O(l)  Mg(OH)2(aq)
2NaOH can also be written: Na+(aq) + 2OH-(aq)
Basic oxides react w/ acids salt and H2O
O2-(s) + 2H+(aq)  H2O(l)
Li2O(s) + 2HCl(aq)  2LiCl(aq) + H2O(l)
MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(l)

32. Chemical Properties: Period 3 Oxides – Acidic Oxides
Acidic oxides – Si, P, S, Cl dissolve in water to produce acidic solutions
Phosphorus(V) oxide reacts with H2O  Phosphoric(V) acid

33. Chemical Properties: Period 3 Oxides – Acidic Oxides

34. Chemical Properties: Period 3 Oxides – Amphoteric Oxides
Aluminum does not affect pH (essentially insoluble)
Will act like a base when reacted with acid
Al2O3(s) + 6H+  2Al3+(aq) + 3H2O(l)
Al2O3(s) + 3H2SO4(aq)  Al2(SO4)3(aq) + 3H2O(l)
Will act like acid when reacted with bases
Al2O3(s) + 3H2O(l) + 2OH-(aq)  Al(OH)4-(aq)

35. Period 3 Oxides Trends

36. You need to know the following:
The equations for
Na2O
MgO
Al2O3 amphoteric property
P4O10
SO3