The Periodic Table

1. The Periodic Table • Dmitri Mendeleev • designed periodic table in which the elements were arranged in order of increasing atomic mass • Henry Moseley • designed periodic table in which the elements were arranged in order of increasing atomic number

2. Periodic law states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically. • We will look in more detail at three periodic properties: atomic radius, ionization energy, and electron affinity.

3. Trends in the Periodic Table The properties of the elements exhibit trends. These trends can be predicted using the periodic table and can be explained and understood by analyzing the electron configurations of the elements.

4. There are two important trends. First, electrons are added one at a time moving from left to right across a period. As this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction, so the electrons become closer to the nucleus and more tightly bound to it.

5. Second, moving down a column in the periodic table, the outermost electrons become less tightly bound to the nucleus. This happens because the number of filled principal energy levels (which shield the outermost electrons from attraction to the nucleus) increases downward within each group.

6. These 2 trends explain the periodicity observed in the elemental properties of atomic radius, ionization energy, electron affinity, and electronegativity.

7. Periodicity of atomic radii • Atomic radii increase down a group • Li  Cs; 2s  7s • Atomic radii decrease going across a period – the • effective nuclear charge , Zeff increases • - a proton is added to the nucleus and • shielding remains constant • Zeff = Zactual – electron shielding

9. The nuclear charge felt by an electron in an outer • shell is called the effective nuclear charge

10. Atomic radii of main-group and transition elements Opposing forces: Changes in n and changes in Zeff Overall Trends (A) n dominates within a group; atomic radius generally increases in a group from top to bottom (B) Zeff dominates within a period; atomic radius generally decreases in a period from left to right

11. Periodicity of atomic radius Large size shifts when moving from one period to the next

13. PROBLEM: Using only the periodic table, rank each set of main group elements in order of decreasing atomic size. Ranking Elements by Atomic Size (a) Ca, Mg, Sr (b) K, Ga, Ca (c) Br, Rb, Kr (d) Sr, Ca, Rb PLAN: Size increases down a group; size decreases across a period. SOLUTION: (a) Sr > Ca > Mg These elements are in Group 2A. (b) K > Ca > Ga These elements are in Period 4. (c) Rb > Br > Kr Rb has a higher energy level and is far to the left. Br is to the left of Kr. (d) Rb > Sr > Ca Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr.

14. IONIZATION ENERGY • Cations are formed when an atom loses one or more • electrons • Na  Na+ + e- • Mg  Mg2+ + 2e- • Energy input is required for this process • Thefirst ionization energy,Ei1, is the minimum • amount of energy required to remove the outermost • electron from an isolated gaseous atom • H + 1312 kJ  H+ + e-

15. In some cases a second and even a third electron • may be removed • Ca + 590 kJ  Ca+ + e- Ei1 • Ca+ + 1145 kJ  Ca2+ + e-Ei2 • Al  Al3+ Ei1, Ei2, Ei3 • For a given element Ei1< Ei2 < Ei3 • - because it is much harder to remove an • electron from a positively charged ion than from • the corresponding neutral atom

16. Compare Ei2 vs. Ei3 for Mg • Mg+ Mg2+ + e- 1451 kJ • Mg2+  Mg3+ + e- 7733 kJ

18. Ionization energy shows a clear periodic trend • Eidecreasesas a group is descended • e.g. Ei for Li > Na > K > Rb > Cs • - the electron is lost from successively higher • energy levels which are further away from the • nucleus

19. There is a gradualincreasein Ei as a period is • traversed • Na < Si < Cl • Part of the reason: atomic radii decrease making the • outermost electrons closer to the nucleus and thus • harder to remove • The increase across the period is not smooth – • breaks occur at Be/B and N/O • Ei for B < Be • Be: 1s2 2s2 Be+: 1s2 2s1

20. Be: 1s2 2s2 Be+: 1s2 2s1 • B: 1s2 2s2 2p1 B+: 1s2 2s2 • In Be, to form Be+ a filled shell is being broken – • this is very energy expensive • On the other hand, B+ has a filled shell; in addition it is easy to remove the single 2p electron • In the next case: • N: 1s2 2s2 2p3 N+: 1s2 2s2 2p2 • O: 1s2 2s2 2p4 O+: 1s2 2s2 2p3 The first three ionization energies of beryllium (in MJ/mol)

23. First ionization energies of the main-group elements Increase within a period and decrease within a group

25. PROBLEM: Using the periodic table, rank the elements in each of the following sets in order of decreasing IE1: Ranking Elements by First Ionization Energy (a) Kr, He, Ar (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs PLAN: IE decreases down in a group; IE increases across a period. SOLUTION: (a) He > Ar > Kr Group 8A elements- IE decreases down a group. (b) Te > Sb > Sn Period 5 elements - IE increases across a period. (c) Ca > K > Rb Ca is to the right of K; Rb is below K. (d) Xe > I > Cs I is to the left of Xe; Cs is further to the left and down one period.

26. PROBLEM: Name the Period 3 element with the following ionization energies (in kJ/mol) and write its electron configuration: IE1 IE2 IE3 IE4 IE5 IE6 1012 1903 2910 4956 6278 22,230 Identifying an Element from Successive Ionization Energies PLAN: Look for a large increase in energy that indicates that all of the valence electrons have been removed. SOLUTION: The largest increase occurs at IE6, that is, after the 5th valence electron has been removed. The element must have five valence electrons with a valence configuration of 3s23p3, The element must be phosphorus. P (Z = 15). The complete electronic configuration is: 1s22s22p63s23p3.

27. ELECTRON AFFINITY • Anions are formed by an atom accepting electron(s) • This process is also accompanied by an energy • change • Theelectron affinity,Eea, is the energy change that • occurs when an electron is added to an isolated • gaseous atom • The energy change is usually negative – the more • negative the Eea, the greater the tendency to form • anions

28. Be + e- Be- Eea = 241 kJ mol-1 • Cl + e-  Cl- Eea = -348 kJ mol- •  Cl form anions easier than Be • The periodic trend is Eea becomes more negative • across a period – trend is not regular • Again there are breaks at Groups 2A and 5A

29. It is very difficult to add an electron to a 2A metal • because its outer 2s orbital is filled • Values for 5A elements are less negative than • expected because they apply to addition of an • electron to a relatively stable half-filled

30. Electron affinities of the main-group elements Negative values = energy is released when the ion forms Positive values = energy is absorbed to form the anion

32. Lets tie it all together:

33. Depicting ionic radii

34. Periodicity of ionic radii • For cations: ionic radii is always less than atomic • radii • - so Li+ < Li • 1s21s2 2s1 • For Li+; 3 p  2 e-  greater attraction here • Li; 3 p  3 e- • For anions: anions are biggerthan their parent • atoms • Cl: 1s2 2s2 2p6 3s2 3p5Cl-: 1s2 2s2 2p6 3s2 3p6

35. Ionic vs atomic radius Ionic size increases down a group Trends in periods are complex For atoms that form more than one cation: the greater the ionic charge, the smaller the ionic radius

37. PROBLEM: Rank each set of ions in order of decreasing size, and explain your ranking: Ranking Ions by Size (a) Ca2+, Sr2+, Mg2+ (b) K+, S2-, Cl- (c) Au+, Au3+ PLAN: Compare positions in the periodic table, formation of positive and negative ions and changes in size due to gain or loss of electrons. SOLUTION: These are members of the same Group (2A) and therefore decrease in size going up the group. (a) Sr2+ > Ca2+> Mg2+ These ions are isoelectronic; S2- has the smallest Zeff and therefore is the largest while K+ is a cation with a large Zeff and is the smallest. (b) S2-> Cl-> K+ (c) Au+> Au3+ The higher the positive charge, the smaller the ion.