Download 1 CHAPTER 5 – THE PERIODIC LAW What types of useful

CHAPTER 5 – THE PERIODIC LAW
What types of useful information can you find on the Periodic Table?
I. History of the Periodic Table
A. Before the Periodic Table was invented, about 63 elements were known. However, they were not
organized and only random properties were known about each of the elements. Scientist (who are
always looking for patterns) wanted to organize these.
B. Dmitri Mendeleev – he made cards for all 63 known elements and included some of their properties.
He tacked them up on his laboratory wall and searched for patterns.
 Mendeleev arranged the elements by increasing atomic mass. When he did this, he found a
periodic pattern of chemical properties.
 In a few cases, he had to put elements out of atomic mass order so their properties would
line up with other elements of similar properties.
 Mendeleev’s Questions:
1. Why are some elements out of “atomic mass order”?
2. What causes this periodicity of chemical properties?
Mendeleev lost the Nobel Prize3 in the year before he died by one vote to the discoverer of
fluorine. Although his table is now the most valuable tool in chemistry, he never won the Nobel
Prize for it.
C. Answer to Mendeleev’s 1st Q:
In 1911, Henry Moseley (who worked with Rutherford) discovered that the positive charge of
the nucleus increases by one unit as you move from one element to the next across
Mendeleev’s table. He decided that the amount of positive charge in the nucleus was based on
the number of protons, and named this the atomic number.
The Periodic Law - The physical and chemical properties of the elements are periodic functions of their
atomic numbers.
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D. The Modern Periodic Table - elements are arranged by increasing atomic number, so that elements
with similar properties are in the same column.
1894 – 1900 – discovery of the noble gases (group 18)
Early 1900’s – lanthanides (rare-earth elements) Period 6, #58 – 71)
Early 1900’s – actinides (al radioactive) Period 7 #90 – 103
 The answer to Mendeleev’s 2nd Q: The reason for periodicity is the arrangement of
electrons around the nucleus of each atom.
Nuts and Bolts Lab
Mendeleev’s P.T. Puzzle Lab
II. Electron Configuration and the Periodic Table
Outer shell e- configurations can be used to explain the chemical properties of the elements.
A. Period numbers and e- configs:
Ex: Find the e- config. of:
H
1s1
He
1s2
Li
Be
B
1s22s1
1s22s2
1s22s22p1
Ne
1s22s22p6
outer level: n=1
outer level: n=2
Pattern?
**Period number (1 -7) = the outer most occupied energy level in the ground state of that atom.
So, what is the outermost occupied level in a K atom? ____4______
a Fr atom? ___7_______
a P atom? ____3______
What period would an atom with the e- config. 1s22s22p63s1? _____3______
How about one with the config. 1s22s22p63s23p64s23d8? ______4_____
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B. group number and e- configuration
1. s-block elements –
Look at Li
1s22s1
Na
1s22s22p63s1
K
1s22s22p63s23p64s1
___1__ outer shell e- in an ___s__ orbital.
Be
1s22s2
Mg
1s22s22p63s2
Ca
1s22s22p63s23p64s2
___2__ outer shell e- in an ___s__ orbital.
Notice any pattern?
** s-block elements: last electron goes into an s orbital.
Also, the group number is the same as the number of outer shell e-.
2. p-block elements – (Groups 13 – 18) have their last e- going into a “p orbital”
Ex:
B
1s22s22p1
__3___ outer shell e-  Group __13___
C
1s22s22p2
__4___ outer shell e-  Group __14___
Ne
1s22s22p6
___8__ outer shell e-  Group __18___
**Rule for s and p block elements
(Groups 1 -2, 13 – 18, called the “main group” elements
 The last digit of the group # tells the # of outer shell e-.
So, how many outer shell e- does Al have? __3___
Se
6 _____
Br
7 _____
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Draw the electron dot notation for S
___________
For Fr
___________
For Rn
__________
How many outer shell e- are in a Bi atom? __5___
What energy level (n) are they in?
__6___
 P-block elements vary in properties within the group, since the “line of demarcation” cuts
through them.
Groups 1: “the alkali metals” (Li, Na, K, Rb, Cs, Fr)
--Highly reactive
-- not found as free (uncombined) elements in nature; always found in compounds
-- all react strongly with water and air, so they are stored under kerosene.
-- all can be cut with a knife
Hydrogen – makes up 76% of universe. Placed above Group 1 (not init, even though it has 1 outer e-)
because
a. It is not a metal
b. It is chemically unlike the Group 1 elements
Group 2: “the alkaline-earth metals” (Be, Mg, Ca, Sr, Ba, Ra)
-- less reactive than Group1, but still founds only as compounds in nature.
Helium - even though it has 2 outer shell e-, it is placed in Group 18 rather than Group 2, because it is
nonreactive like the other noble gases.
Group 17 – the “halogens” (F, Cl, Br, and I)
Group 18 – the “noble gases” (He, Ne, Ar, Kr, Xe, Rn)
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3. D-block elements – (Groups 3 -12)
-- also called the “transition elements”
-- e- configuration is somewhat less predictable for these elements (although most are normal)
-- all have typical metallic properties
-- generally have 2 valence e- (outer shell)
4. f-block elements – (belong between groups 3 and 4; placed below the rest of the P.T.) usually have 2
valence e-
Pass out lrg. P.T.
--Lanthanides – found in Pd. 6 – also called “rare-earth elements” although they aren’t rare.
Shiny, reactive metals.
-- Actinides – found in Pd. 7. Unstable and radioactive. The 1st four are found naturally; the rest
have only been made in the lab.
QUIZ
III. Electron Configuration and Other Periodic Properties
A. Atomic Radius – distance from the nucleus to the edge of an atom.
As atomic # increases across a period, more e- are added to the same shell, but more p are
also added to the nucleus. This stronger + charge in the nucleus pulls the e- closer,
decreasing the radius.
As you go down a group, new e- are added to a new shell, farther from the nucleus, so radius
increases.
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B. Ionic Radius - Ionic sizes relative to other ions. Follows the same trends as atomic radius.
Cation – (cat-ion) a positively charged ion
-- Formed by the loss of 1 or more e-- Causes a decrease in radius compared to the neutral atom
-- Most metals form cations (left side of table; low # of valence e-, easier to lose a few to
get to octet than to gain several).
Anion – (an-ion) – a negatively charged ion
-- Formed by gaining 1 or more e-- causes and increase in radius compared to the neutral atom
-- Most nonmetals form anion (right side of table; high # of valence e-, easier to gain a
few to get octet than lose several).
C.
D. Electronegativity – the ability of an atom within a chemical compound dto attract
electrons.
-- Based on whether attracting e- gets them closer to an octet, and based on size
(smaller atoms attract e- more because the e- are closer to the nucleus).
Fluorine – the most electronegative element. Assigned an arbitrary EN of 4.0 (all other
elements have a lower EN)
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In the compound HF, to which atom will the e- be closer? F
How about the compound NaCl? Cl
The compound H2O? O
D. The Electron affinity - of an atom is the energy change when an electron is added
to the neutral atom to form a negative ion. This property can only be measured in an atom in
gaseous state.
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E. Ionization Energy - the minimum energy required to remove (to infinity) an electron from the atom
or molecule
Summary of Periodic Trends
Periodic Law Puzzle
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