1. Periodic Table – Organizing the Elements
Chapter 5.4
& Chapter 14

2. Dmitri Mendeleev About 70 elements had been found by the mid 1800’s
Mendeleev was the first to organize them in a systematic way

3. He listed the elements in order of increasing atomic mass
Arranged the elements in columns so those with similar properties were side by side

4. He left blank spaces where nothing fit
He predicted the physical properties of the missing elements
He was mostly correct

6. Henry Moseley
Moseley determined the atomic # of the elements and arranged the table by atomic number instead of atomic mass

7. The modern periodic table is arranged by atomic number
The periodic table has atomic # increasing from left to right & top to bottom

8. Metals Vs Nonmetals
The table can be divided into 3 main types of elements – metals, nonmetals & metalloids
Metals reside on the left side of the table
Nonmetals are on the upper right side of the table
Metalloids are b/tw the metals & nonmetals

9. Metals are typically shiny (luster) & are good conductors of heat and electricity
Metals are all solids, except Mercury (Hg) which is a liquid
Nonmetals are usually dull (no luster) & are poor conductors of heat and electricity
Nonmetals can be solid (C & S), liquid (Br) or gas (O & H)

10. Metalloids
Metalloids are their own category of elements – there are only 7 metalloids!
B, Si, Ge, As, Sb, Te, & At – they border the stair-step line that most periodic tables include – but not all.
Metalloids have properties in b/tw the metals & nonmetals
Si is used in computers b/c it is a semi-conductor

11. Periodic Law
The horizontal rows on the periodic table are called periods
Properties change as you move across a period

12. The properties repeat when you move from one period to the next
Periodic Law: there is a periodic repetition of the chemical & physical properties of the elements

13. Groups
Each vertical column is called a group or family
Elements in the same group have similar properties

14. Groups have a number and a letter (pg 124)
The group with Li, Na, K etc is called Group 1A
Group 1A elements are called the alkali metals
Alkali metals are the MOST active metals

15. Group 2A are the Alkaline Earth Metals
Alkaline Earth Metals are not as active as the alkali metals, but react with many substances

16. Group 7A are called the Halogens (F, Cl, Br, I)
Halogens are the most active nonmetals
Group 0 (8A) are the noble gases (He, Ne, etc)
The noble gases are mostly UNREACTIVE!

17. Groups 3A to 6A
Groups 3A to 6A do not have specific names – they are usually named by the top element of the group.
For instance, The Carbon family is Group 4A.
These groups can contain metals, metalloids and nonmetals!

18. All group A elements are called the representative elements
They exhibit a wide range of physical & chemical properties

19. Group B elements are the transition & inner-transition metals
Gold & silver are transition metals
Uranium is an inner-transition metal

20. The elements can also be classified by their electron configuration
Electrons play the most important part in determining the properties of elements

21. Write the electron configurations for the Alkali Metals
What similarities do you see?
The Halogens?
The Noble Gases?

22. The noble gases have their outermost s & p sublevels filled completely
The Representative Elements have their outermost s & p sublevels partially filled

23. The Transition Metals – their outermost s & nearby d sublevels contain electrons
The Inner Transition Metals – their outermost s & nearby f sublevels contain electrons

24. s block, p block, d block & f block Where are they?
Blocks
The Table can be broken up into blocks - tell you the outermost sublevels that are filled
s block, p block, d block & f block
Where are they?

26. “S” block Groups 1 & 2 Alkali Metals and Alkaline Earth Metals
Electron Configuration ends in an S Sub-level.

27. “P” block Groups 13 thru 18 (or 3A - 8A)
Includes halogens and noble gases (except He)
Electron Configuration ends in a P Sub-level.

28. “D” block Groups 3 thru 12 Transition Metals Au, Ag, Fe, Pt, etc
Electron Configuration ends in a D Sub-level.

29. “F” block “Inner Transition Metals” Includes Uranium
Electron Configuration ends in an F Sub-level.

30. Periodic Trends
An element’s placement in the periodic table determines characteristics like the size of the atom, its ability to attract electrons and the stability of its electron configuration.

31. Atomic Radius Size of atoms of each element:
How will the size of atoms change as we proceed down a group?
i.e. Compare the sizes of Li and Na.
From Li to Na, we add an entire energy level, therefore the size increases.

32. How will the size of atoms change as we proceed across a period?
Compare C, N and O. Which is largest?
Oxygen has the most electrons.
However, it also has the most protons.
The outermost electrons of Oxygen are in the same sub-level as C and N.

33. Oxygen’s greater nuclear charge attracts the electrons, causing the atom to contract!
Oxygen is the smallest of the three, Carbon is the largest.
Atomic Radius decreases as we go across a period from left to right and increases going down a group.

35. Rank the following sets in order of decreasing Radius.
Examples
Rank the following sets in order of decreasing Radius.
S, Cr, Se, Sr, Ne
Sr, Cr, Se, S, Ne
Fe, N, Ba, Ag, Be
Ba, Ag, Fe, Be, N

36. Ionization Energy
Amount of energy required to remove a valence electron from an atom.
The more stable an element is, the harder it will be (more energy is required) to remove an electron.
Some elements become more stable by losing an electron so they lose electrons easily (less energy needed).

37. How does ionization energy vary within a group (compare Li and Na)?
The electron to be removed from Na is further from the nucleus than Lithium’s electron.
Sodium’s electron is held more loosely and therefore easier (less energy) to remove.

38. How does ionization energy vary across a period
How does ionization energy vary across a period? (Compare elements in 3rd period)
Sodium attains a Noble Gas configuration by losing an electron, so little energy is required.
Magnesium is somewhat stable due to a full 3s sub-level, so more energy is needed.

39. Ionization energy increases across a period & decreases down a group.
Argon is a Noble Gas. Due to its stability, it is very difficult (much energy needed) to remove an electron.
Chlorine has no stability in its configuration, so it is easier to remove an electron.
Ionization energy increases across a period & decreases down a group.

41. Rank the following sets in order of decreasing Ionization Energy.
Examples
Rank the following sets in order of decreasing Ionization Energy.
K, Zn, Cs, Ar, P
Ar, P, Zn, K, Cs
C, He, Ag, Pt, Sn
He, C, Sn, Ag, Pt

42. 2nd Ionization Energy
After an element loses one electron, it may lose another.
Sometimes it is easier to lose the second electron than the first.
Sometimes the first electron is easier to remove
This depends on the stability of the electron configuration after it loses the 1st electron

43. 2nd Ionization Energy
Magnesium will be more unstable after losing 1 electron, so the 2nd electron will be lost more easily
Sodium becomes very stable after losing just 1 electron, so it will be more difficult to lose the second electron

44. Electronegativity
Describes an element’s attraction for an electron in a covalent bond.
Elements that need electrons to complete an energy-level will have a high electronegativity.
Elements that want to lose electrons have low electronegativities.

45. How does Electronegativity vary within a group? (compare F and Cl)
Both elements need an electron to complete a p sub-level.
Fluorine’s p sub-level is closer to its nucleus, so it has a greater magnetic attraction for a free electron.
F has a higher electronegativity!

46. How does electronegativity vary across a period? (period 2)
Fluorine benefits the most by gaining an electron, so it has the highest electronegativity.
Lithium, which wants to lose an electron has very little attraction for an additional electron.

47. Carbon can gain electrons but sometimes loses them as well, so its electronegativity is between F and Li.
Noble Gases have no Electronegativity!
Electronegativity increases across a period and decreases down a group (Noble Gases omitted).

49. Rank the following sets in order of decreasing Electronegativity:
Examples
Rank the following sets in order of decreasing Electronegativity:
Cu, F, Mn, Sr, Si
F, Si, Cu, Mn, Sr
Al, Ca, S, Cl, Fe
Cl, S, Al, Fe, Ca

50. Atomic Radius vs. Ionic Radius
When atoms become ions, their size will change:
Metals will lose electrons and become smaller than the neutral atom (Na+ is smaller than Na)
Nonmetals will gain electrons and become larger than the neutral atom (F- is larger than F)