1. Beyond protons, neutrons, and electrons
The Periodic Table
Beyond protons, neutrons, and electrons

2. It wasn’t always like this…

3. Early PT Folks Johann Dobereiner JAR Newlands (1864) Law of Octaves
Johann Dobereiner
Triads- groups of 3 with similarities/ trends
Cl, Br, I – the properties of Br were intermediate to those of Cl and I
Limited to some groups, not effective with others
JAR Newlands (1864) Law of Octaves
Every eight elements the pattern repeats itself, similar to a musical scale repeating every 8 notes
Not generally well received; people thought him a fool

5. The Modern Periodic Table
The original PT was arranged by mass
By Dmitri Mendeleev and J Lothar Meyer in 1869
Mendeleev predicted the existence of unknown elements (which turned out to be Ge, Sc, and Ga), and predicted their properties from the patterns he saw
Mendeleev corrected the assumed atomic masses for elements (In, Be, U)
These are reasons why he is credited with the first periodic table and is dubbed “The Father of the Modern Periodic Table” over Meyer

6. Ekasilicon

7. Changes….
Henry Mosley changed the table to be organized by atomic number (Z) instead; it then more closely followed trends/ patterns

8. Patterns (Periods) and the PT
We see patterns for many things, including
Atomic number *(not a periodic pattern, but a pattern)
Electron configuration
Atomic radii
Ionization energy
Electron affinity
Electronegativity
Activity
Density

9. The Periodic Law
Mendeleev says "The properties of the elements are a periodic function of their atomic masses"
We now say: “When atoms are arranged by increasing atomic number, the physical and chemical properties show a (repeating) pattern”
Hence, we call the table of elements the PERIODIC table (go figure)

10. The Modern Periodic Table
Consists of boxes containing:
Element name and symbol
Atomic number and atomic mass
Groups:
Columns (up and down)
18 of them
Periods:
Rows (left to right)
7 of them
Representative elements:
elements in groups 1,2 and 13-18
Have a wide range of chemical and physical properties

11. Metals Conductors of heat and electricity
Make cations (lose e- to become + charged)
Malleable (made into sheets)
Ductile (made into wire)

12. Nonmetals Are a brittle solid or a gas
Make anions (gain e- to become - charged)
Covalently bond to each
other

13. Semi-metals (AKA Metalloids)
Characteristics of both metals and nonmetals
More metallic as you go down PT

14. Valence electrons and the PT
PT also shows trends in valence electrons. Elements in:
Column 1 have 1 valence e-
Column 2 have 2 valence e-’s
Columns 13 to 18 have valence e-’s equaling the column # - 10
Column 13 have 3 valence e-’s (13- 10)
Column 14 have 4 valence e-’s (14- 10)
Column 15 have 5 valence e-’s ( )

15. Periods, e- configuration and the PT
Period trends are seen in the electron configuration
For columns 1,2 and 13-18, the period on the table matches the energy level (ring, n value)
Alkali metals are #s1 (# is period)
Alkaline earth metals are #s2 (# is period)
Halogens #p5 (# is period)
Noble gases #p6 (# is period)
Different trend for d and f blocks
Transition metals #d (# is period -1)
Inner transition metals are #f (# is period -2)

16. e- configuration and the PT
Group trends are also seen in the electron configurations
Columns 1 and 2 end in s (s1 and s2 )
Columns 13 to 18 end in p (p1 to p6 )
Transition metals end in d (d 1 to d 10 )
Inner transition metals end in f (f 1 to f 14 )

17. Blocks and l* * * orbital shape
Blocks and l* * * orbital shape
The blocks you already know correspond to the orbital of the last (outermost) e- , or valence e-s occupied

18. Octet Rule
“Atoms gain, lose, or share electrons in order to create a full outer shell”
This is typically going to be eight electrons
H and He are exceptions; wanting to fill the 1s orbital
H gains an electron to become H- , with the same electron configuration as He
H may want to go to no electrons, which is considered “full” even though it is empty
H+ and He+2 would have no electrons left
The law can be used to predict several properties

19. Common Ions of Elements+1 +3
+/-4 -3 -1 -2 +2
Variable, always +

20. Effective Nuclear Charge
Nuclear charge – the attraction felt for an electron by the nucleus
Electrons are both
attracted to the nucleus and
repelled by other electrons.
The nuclear charge that an electron experiences depends on both factors.
This effects all periodic properties

21. Figure 7.4 Variations in effective nuclear charge for period 2 and period 3 elements.

22. Atomic Radii Half the distance between adjacent nuclei
Half the distance between adjacent nuclei
½ (2R)= atomic radius

23. Atomic Radii The radius increases as you go down a group
The radius increases as you go down a group
This is because n increases
The radius decreases as you go across a period
(Yes, this is counterintuitive)
Due to the fact that you add e- as you add p+, so the nucleus is more positively charged, and each electron has the same negative charge
Results in each electron being more attracted to the (increasingly) more positive nucleus, and being pulled in closer
Sort of like making a magnet more powerful- it will decrease the distance where it will pull objects towards it

25. Ionic Radii Larger than the neutral atom Cations (+) Anions (-)
Cations (+)
Smaller than the neutral atom
The electrons have less repulsion, and pull in closer to the nucleus
Anions (-)
Larger than the neutral atom
More electrons = more repulsion = larger electron cloud

27. Ionization Energy (Heretofore called IE)
IE is the amount of energy needed to remove an electron from an atom
(specifically, an isolated atom of the element in the gas phase)
Measure in kJ/ mol
Al(g)Al(g)+ + e I1 = 580 kJ/mol
Al(g)+ Al(g)+2 + e I2 = 1815 kJ/mol

28. IE, continued
The Energy needed to remove the first electron from an element is the 1st IE
The Energy needed to remove the second electron is known as the 2nd IE

29. Successive IE
There are also 3rd, 4th, 5th , and so on IEs (which are successive IEs), until you can’t pull any more off
It takes more energy to remove successive electrons than to remove the first
Due to the fact that there are then more protons than electrons, and the stronger positive charge will then act on the remaining electrons to hold them to the atom
(Remember that the charge on the nucleus increases while the charge on each electron remains the same, causing more pull by the nucleus on each individual electron)

30. Why IE?
Since electrons (-) want to hang around the atom (due to the + protons in the nucleus pulling on them), it takes energy to remove electrons
In general
The smaller that atom, the more energy it takes to remove an electron
Because the electron is closer to the nucleus than in a larger atom
The fewer electrons that atom possess, the harder it is to remove an electron
Because it will hang on to them tighter as they are closer to the + charged nucleus;
also, the less repulsion between electrons

32. 1st IE

33. Things to keep in mind…
Remember (from coming up with the abbreviated electron configurations) that:
Inner core electrons are those electrons from previous Noble Gas
Valence electrons are the electrons that are on the exterior of an atom
These are the electrons that are responsible for the behavior (properties) of the element

34. Successive IEs Are higher than the first
Due to the fact that there is going to be more protons than electrons at that point, resulting in a stronger attraction on the remaining electrons than there was in the first place
Basically increasingly larger jumps as each electron is removed
One jump is usually much larger than the others, because once the inner core configuration is reached, electrons are removed from the inner core, taking a lot more energy
Much bigger difference between positive nucleus and negative electron

35. Successive IEs I1 I2 I3 I4 I5 I6 I7 495 4560 735 1445 7730 580 1815Na
495
4560 Mg
735
1445
7730 Al
580
1815
2740
11600 Si
780
1575
3220
4350
16100 P
1060
1890
2905
4950
6270
21200
1005
2260
3375
4565
6950
8490
27000 Cl
1255
2295
3850
5160
6560
9360
11000 Ar
1527
2665
3945
5770
7230
8780
12000

37. Electronegativity (Eneg)
The ability of an atom to attract electrons in a bond
Some atoms share electrons easily, others are electron hogs
The ability to share is rated (usually) from 0 to 4
Elements with 0 Eneg share easily
Elements with a high (close to 4) Eneg don’t share e- well

38. Electronegativity Trends
If it normally goes +, it has a low Eneg
If it normally goes -, it is has a high Eneg
The smaller it is, the higher the Eneg
The larger it is, the lower the Eneg
Noble gases, which normally take no charge, we say have no Eneg values

39. Electronegativity Trends

40. Metallic character
Metallic character is acting like a metal (conductive, shiny, malleable,etc)
All elements possess from very low to very high metallic character.
The scale is from Fr to F.
Fr has the most metallic character and F has the least.
In groups,  metallic character increases with atomic number because each successive element gets closest to Fr.
In periods,  metallic character decreases when atomic number increases because each successive element gets closest to F.

41. Reactivity
The nature (metal, non-metal, semi-metal) makes a difference in how an element’s chemical reactivity
The trends are characterized by their nature

42. Metals reactivity trend
In groups, reactivity of metals increases with atomic number because the ionization energy decreases.
In periods, reactivity of metals decreases when atomic number increases because the ionization energy increases.

43. Nonmetals reactivity trend
In groups, reactivity of non-metals decreases when atomic number increases
because the electronegativity decreases
Relate to size- it increases.
In periods, reactivity of non-metals  increases with atomic number
because the electronegativity increases.
Relate to size- radii decreases
Remember, the radii would have an effect on this

44. Density: in general Density of solids is greatest Density of gases
Measured in g/cm3
Highest in center of table (d- block)
Density of gases
Measured in g/L at Standard Temp &Pressure (STP, which is 1atm and 0°C)
Increases as you go down a group
Decreases as you go across the table
Density of liquids
Measured in g/mL
Density of Hg is greater than that of Br2

45. Density

46. Density v atomic number

47. Summing it up (again)

48. Summary chart again