1. CHE 106 Chapter 7: Periodic Table

3. Development of the Periodic Table
The Periodic Table organizes information about elements and allows us to recognize trends in reactivity and behavior of elements. The trends that we will discuss in this chapter can all be related back to the electronic arrangement.

4. Development of the Periodic Table
Most elements are not found in nature in their elemental form, instead forming compounds to achieve stability. Prior to the 20th century, many elements weren’t known because of a lack of technology. Prior to 1800, only 31 elements were known, they were found in nature in their elemental form. By 1865, there were 65.

5. Development of the Periodic Table
1869: Dmitri Mendeleev: recognized chemical and physical properties that repeated periodically if the elements were arranged in increasing atomic weights. He left blank spaces to leave room for elements that had yet to be discovered, as he arranged elements into groups based on their reactivity and behavior.

6. Development of the Periodic Table

7. Development of the Periodic Table
Moseley: Discovered atomic number concept using x- rays. Rearranged the periodic table by order of increasing atomic number… clarified some of the inconsistencies. Periodic Table: Arranged into rows and columns: -Periods: Same principal quantum number (n), same “shell” -Groups: Increasing value of n

8. Nuclear Charge and Shielding
Two forces working on an electron: attractive force from the nucleus and repulsive force for nearby attraction. The outermost electrons behavior depends largely on the balance between these two forces. When there are many electrons in an atom, the Zeff decreases because of the increased amount of shielding. Zeff = Z – S

9. Nuclear Charge and Shielding
Recall from Chapter 6: When Zeff decreases, the electrons have more energy. And the level of shielding within the same principal quantum number: s>p>d>f.

10. Nuclear Charge and Shielding
As a general trend: as we move from left to right, the effective nuclear charge increases.
- The core electrons who are responsible for shielding remains the same within a period, therefore S is constant, while Z increases with atomic number.
As we go down a group, the effective nuclear charge changes very little, increases but only slightly.
- As you increase nuclear charge going down, you also increase the shielding value.

11. Atomic Sizes
As the definition of the atom changes from a hard sphere to the quantum mechanical model – the boundaries of where the atom started and stopped was impossible to define. We estimate the atomic radii by assuming that when atoms bond together, they “touch” and we can approximate that distance.

12. Atomic Sizes
Bonding atomic radius: half the nucleus to nucleus distance when two atoms are bonding to one another. Nonbonding Atomic Radius: the shortest distance separating two nuclei during a collision.

13. Atomic Sizes
Periodic Trends:
Radii increase from top to bottom within a group.
You are increasing the principal quantum number, outer electrons are farther from the nucleus.
2. Radii decrease from left to right across a period.
Due the the increasing Zeff, the pull inward is stronger, shortening the radius.

14. Atomic Sizes
Li < Na < K < Rb < Cs, due to increasing n values.
Li > F because of Zeff changes. Rb K
Transition
Metals Na
Transition
Metals
Atomic
Radius
(Å) Li I F Cl Br
Atomic Number

15. Atomic Sizes

16. Atomic Sizes: Ionic Radius
Cations: Smaller than their parent atoms
You have reduced the strength of electron-electron repulsions and this allows the effective nuclear charge to have a stronger pull on the remaining electrons.
Anions: Larger than their parent atoms
- By adding electrons, there are more repulsive forces, so this stretches the atom and it increases in size.

17. Atomic Sizes: Charge Effects
(232 pm)
Fe+2
(152 pm)
Fe+3
(128 pm) F
(128 pm)
F -1
(272 pm) Li
(304 pm)
Li+
(120 pm)

18. Atomic Sizes: Ionic Radius
Other trends: Within a group, all of the ions will have the same charge… As you move down a group, the size of the ion increases. Isoelectronic Series: group of ions with the same number of electrons: When listed in order of increasing atomic number (increasing Zeff), and the same number of electrons, the sizes of the ions decrease.

19. Atomic Sizes: Ionic Radius

20. Atomic Sizes: Ionic Radius

21. Atomic Sizes: Ionic Radius
Example: Arrange the following atoms in order of increasing atomic radius: Na, Be, Mg ANSWER: Be < Mg < Na

22. Ionization Energy E(g)  E+1(g) + 1 e-
Electrons are exchanged during chemical reactions. Ionization Energy (IE) measures how strongly an atom holds on to its electrons. Also can be defined as the minimum energy necessary to remove an electron from the ground state of an isolated gaseous atom. 1st IE: removing 1st electron 2nd IE: removing 2nd electron: takes more energy 3rd IE: removing 3rd electron: takes even more energy
E(g)  E+1(g) e-

23. Ionization Energy
Within a period: I1 increases with increasing atomic number due to increasing Zeff. Within each group: I1 decreases with increasing atomic number due to increasing atomic radius.

24. Ionization Energy
Which element: Na, Ca, S would have the highest 2nd ionization energy? ANSWER: Na  removing an inner electron requires a tremendous increase in energy than removing an outer electron.

25. Ionization Energies (IE)
1s s p 2p 2p
He (1s2)
(1s22s22p6) Ne
(1s22s22p3) N
F (1s22s22p5)
1st IE
(kJ/mol)
H (1s1)
O (1s22s22p4)
Be (1s22s2)
C (1s22s22p2)
B (1s22s22p1)
Li (1s22s1)
Atomic Number

26. Ionization Energies (IE)
Screening and increasing Zeff He Ne
Zeff = 1 to 2
Zeff = 2 to 3
Zeff = 5 to 6 F
Zeff = 1 to 2
1st IE
(kJ/mol) N H O Be
Zeff = 4 to 5 C B
2s to 2p Li
Zeff = 1 to 2
Atomic Number

27. Ionization Energies (IE)
1s s p 2p 2p He Ne
pairing energy F
1st IE
(kJ/mol) N H O Be C
n = 1 to 2 B
increasing n and
pairing energy Li
Atomic Number

28. Ionization Energy

29. Ionization Energy Elem. Elec. Config. I1 I2 I3 I4 I5 I6
Na [Ne]3s
Mg [Ne]3s
Al [Ne]3s23p
Si [Ne]3s23p
P [Ne]3s23p
S [Ne]3s23p
Core Electrons
Being Removed

30. Ionization Energy
Example: Based on the trends discussed: which of the following atoms, B, Al, C, or Si – has the lowest 1st ionization energy? ANSWER: Al Example: Is it easier to remove a 2p electron from the oxygen atom or a nitrogen atom? ANSWER: It would be easier to remove it from oxgyen.

31. Electron Affinity
Electron Affinity: the energy associated with adding an electron to a gaseous atom. For most atoms and all cations – energy is released when the electron is added. Negative EA value when energy is released upon adding an electron. (Exothermic). Positive EA value when energy is required to add an electron (Endothermic).

32. Electron Affinity

33. Electron Affinity
The greater the attraction between the atom and the electron (greater Zeff) the more negative the EA value becomes. EA values become more negative as you go from left to right across the period. * The Halogens have the most negative EA values * Noble Gases are positive because they have filled shells, to open up a new sublevel – not favorable. *Group 2: EA values are positive because they have to open up the p sublevel, which is energetically unfavorable.* *Group 15: have half filled p sublevel, adding another electron to an orbital increases electron repulsion: EA values are not very negative. EA values don’t change tremendously as you move down a group. Adding electrons to 2p vs 3p vs 4p. Distance from nucleus increases so attraction for an electron decreases as well as reducing the electron- electron repulsion because the orbital are bigger and more spread out as you go farther down.

34. Trends and Group Properties
Trends in radius, IE and EA are helpful to predict chemical behavior and are atomic properties. Sometimes helpful to describe properties of groups of atoms.

35. Trends and Group Properties
Metals: conductors of heat and electricity, shiny luster (many are silvery). Malleable and ductile and most are solids at room temperature except Hg. Tend to form metal oxides that are basic and form cations. Low IE. Non-metals: poor conductors of heat and electricity (insulators) brittle and dull. Most of them are non-solids at room temp. The non-metal oxides form acidic solutions and tend to form anions. High IE. Metalloids: have properties of both. Metalloid oxides are amphoteric.

36. Trends and Group Properties
Metal Oxides are basic: Metal Oxide + Water  Metal Hydroxide MgO + H2O  Mg(OH)2 Na2O + H2O  2 NaOH Metal Oxide + Acid  Salt + Water MgO + 2HCl  MgCl2 + H2O

37. Trends and Group Properties
Non-metal oxides are acidic: Nonmetal oxide + water  acid P4O H2O  4 H3PO4 CO2 + H2O  H2CO3 Nonmetal oxide + base  salt + water CO2 + 2 NaOH  N2CO3 + H2O B2O3 + 6 NaOH  2 NaBO3 + 6 H2O

38. Trends and Group Properties
Alkali Metals: Most reactive metals, soft metallic solids. As you move down the group the following trends: IE decreases, MP/BP decreases, tendency to share electrons (covalency) decreases Radius increases

39. Trends and Group Properties
Alkali Metals
Readily lose 1 electron, never found in nature alone – always in a compound. Use electrolysis of a fused molten salt to isolate.
React with hydrogen to form hydrides – hydrogen assumes a -1 oxidation state.
React with water in very exothermic reactions to form basic solutions. Reactions become more violent as you move down the group.

40. Trends and Group Properties
Alkaline Earth Metals
Trends in IE, radius, mp, bp and covalency mirror the trends found in Group 1 elements. Tend to lost 2 electrons.
Both Group 1 and Group 2 metals are colorless in solution, but give bright colors in flame tests.
Also react with water to form basic solutions, just less violently, more slowly than Group 1 metals.

41. Trends and Group Properties
Group 16: Chalcogens/ Oxygen group
IE decreases and covalency still decreases as you move down the group.
Radius increases and the MP, BP increases.
Oxygen: allotropes are O2 and O3. High tendency to gain electron – “oxidize” other elements. S and O react very similarly.

42. Trends and Group Properties
Group 17: Halogens “Salt formers”
IE decreases and covalency still decreases as you move down the group.
Radius increases and the MP, BP increases.
Many are diatomic in elemental state, highest EA.
Fluorine has highest EA – can take electrons from nearly every element.
They react with hydrogen to form compounds that when dissolved in water to form acids. Most are strong acids.

43. Trends and Group Properties
Group 18: Noble Gases
IE decreases and covalency still decreases as you move down the group.
Radius increases and the MP, BP increases.
Very unreactive – closed shells. Monoatomic.
Xe will react with strong oxidizers, Kr only reacts with F and He, Ne and Ar do not form compounds.