Periodic Table and Periodic Law


The Periodic Table got its name because of the repeating pattern of chemical & physical properties .

Mendeleev ordered his periodic table with elements arranged in order of increasing atomic mass .


Mendeleev noticed there seemed to be a repeating pattern of properties such as densities, formulas with oxygen & hydrogen, boiling or melting points every 8 or 18 elements . He called this repeating quality, periodic. (Periodic ~ according to a pattern)

He started new rows so that the elements having the same properties would be aligned in each column.
Similar properties in each column
Holes for undiscovered elements


He also noticed some gaps missing elements

Based on his periodic table Mendeleev predicted the properties of the missing elements. Others tried to prove him wrong, but it turns out that he was right. Scientists soon found the missing elements and Mendeleev was very close with his predictions.

Atomic mass
Density
Melting point
Color
Obtained in
Will form
Density of oxide
Solubility
Mendeleev’s Predictions
Germanium (Ge) Ekasilicon (Es)
72
5.5 g/cm 3
High
Dark gray metal
K
2
EsF
6
EsO
2
4.7 g/cm 3
Slightly in HCl

Mendeleev’s Predictions
Atomic mass
Density
Melting point
Color
Obtained in
Will form
Density of oxide
Solubility
Ekasilicon (Es)
72
5.5 g/cm 3
High
Dark gray metal
K
2
EsF
6
EsO
2
4.7 g/cm 3
Slightly in HCl
Germanium (Ge)
72.61

Mendeleev’s Predictions
Atomic mass
Density
Melting point
Color
Obtained in
Will form
Density of oxide
Solubility
Ekasilicon (Es)
72
5.5 g/cm 3
High
Dark gray metal
K
2
EsF
6
EsO
2
4.7 g/cm 3
Slightly in HCl
Germanium (Ge)
72.61
5.323 g/cm 3

Mendeleev’s Predictions
Atomic mass
Density
Melting point
Color
Obtained in
Will form
Density of oxide
Solubility
Ekasilicon (Es)
72
5.5 g/cm 3
High
Dark gray metal
K
2
EsF
6
EsO
2
4.7 g/cm 3
Slightly in HCl
Germanium (Ge)
72.61
5.323 g/cm 3
945 o C

Mendeleev’s Predictions
Atomic mass
Density
Melting point
Color
Obtained in
Will form
Density of oxide
Solubility
Ekasilicon (Es)
72
5.5 g/cm 3
High
Dark gray metal
K
2
EsF
6
EsO
2
4.7 g/cm 3
Slightly in HCl
Germanium (Ge)
72.61
5.323 g/cm 3
945 o C
Gray metal

Mendeleev’s Predictions
Atomic mass
Density
Melting point
Color
Obtained in
Will form
Density of oxide
Solubility
Ekasilicon (Es)
72
5.5 g/cm 3
High
Dark gray metal
K
2
EsF
6
EsO
2
4.7 g/cm 3
Slightly in HCl
Germanium (Ge)
72.61
5.323 g/cm 3
945 o C
Gray metal
K
2
GeF
6

Mendeleev’s Predictions
Atomic mass
Density
Melting point
Color
Obtained in
Will form
Density of oxide
Solubility
Ekasilicon (Es)
72
5.5 g/cm 3
High
Dark gray metal
K
2
EsF
6
EsO
2
4.7 g/cm 3
Slightly in HCl
Germanium (Ge)
72.61
5.323 g/cm 3
945 o C
Gray metal
K
2
GeF
6
Forms GeO
2

Mendeleev’s Predictions
Atomic mass
Density
Melting point
Color
Obtained in
Will form
Density of oxide
Solubility
Ekasilicon (Es)
72
5.5 g/cm 3
High
Dark gray metal
K
2
EsF
6
EsO
2
4.7 g/cm 3
Slightly in HCl
Germanium (Ge)
72.61
5.323 g/cm 3
945 o C
Gray metal
K
2
GeF
6
Forms GeO
2
4.70 g/cm 3

Mendeleev’s Predictions
Atomic mass
Density
Melting point
Color
Obtained in
Will form
Density of oxide
Solubility
Ekasilicon (Es)
72
5.5 g/cm
High
3
Dark gray metal
K
2
EsF
6
EsO
2
4.7 g/cm 3
Slightly in HCl https://www.youtube.com/watch?v=kuQ0Um4Wcz0
Germanium (Ge)
72.61
5.323 g/cm 3
945 o C
Gray metal
K
2
GeF
6
Forms GeO
2
4.70 g/cm 3
Not dissolved in HCl


In the modern Periodic Table, elements are arranged in order of increasing atomic
.


The horizontal rows are called

The vertical columns are called

The patterns of properties repeat in each new row, so elements in the
, have similar chemical and physical properties. This is called the periodic


1A 2A 3A 4A 5A 6A 7A 8A


The group B elements (the d block) are called the transition elements.

The group B elements (the f block) are called the inner transition elements.
Transition
Inner
Transition


There are 3 main classes of elements: metals,
nonmetals, metalloids .


The electron structure of an atom determines many of its chemical & physical properties.
For the group A elements, the group number equals the number of valence electrons.

(Except for Helium=2)


The Octet rule states that atoms lose, gain, or share electrons in order to gain a full set of 8 valence electrons.

This noble gas configuration is very stable .
(the exceptions are Hydrogen and helium which will have a stable set of
2 electrons in the 1 st energy level).
Using the octet rule, you can predict which ions will likely form.
Neon has full outer shell : stable & nonreactive


Metals are electron donors . They tend to lose electrons and become + (positive) charged ions.
Called a cation.
Na has 1 valence electron
Must lose 1 electron to have full outer shell
(easier to lose 1 than gain 7)
Creates a ion with a
+1 charge


Nonmetals are electron acceptor . They tend to gain electrons and become – (negative) charged ions. Called an anion.
Chlorine has 7 valence electrons
Need 1 more electron to gain full outer shell
Creates -1 ion


: Tends to decrease across the period. Electrons are being added in the same energy level so increased attraction between the larger number of + protons pulls the – electrons closer.


: Tends to increase down the group, because you are adding energy levels, which shield the valence electrons from the pull of the nucleus.


Losing electrons produce + charged ions, which are smaller than the parent atom.

Na atom
1 valence electron
Effective nuclear charge on remaining electrons increases.
11p
+
Remaining e- are pulled in closer to the nucleus. Ionic size decreases.
Valence e- lost in ion formation
Result: a smaller sodium cation, Na +


When atoms gain electrons, they become – charged ions, which are larger than the parent atom.

Chlorine atom with 7 valence e-
A chloride ion is produced. It is larger than the original atom.
17p
+
One e- is added to the outer shell.
Effective nuclear charge is reduced and the e- cloud expands.


As with the atomic radius size, decrease across the periods an increase down the group.

Notice difference between cations & anions; but trends are still same decreases increases

Ionic Radius – size as compared to the
element
- Positive cations are always smaller than neutral atom (removal of e ) often remaining e are in a lower energy level.

Na 1s 2 2s 2 2p 6 3s 1 Na 1+ 1s 2 2s 2 2p 6
Anions get bigger than neutral atom. Adding e creates more e repulsion.

Cl 1s 2 2s 2 2p 6 3s 2 3p 5 Cl 11s 2 2s 2 2p 6 3s 2 3p 6


Is the amount of energy required to pull the 1 st valence electron away from the atom.

Elements with high ionization energy are
[unlikely, likely] to lose electrons. Atoms with low ionization energy [easily, don’t really] lose electrons


Generally metals have low ionization energies and easily form positive ions… nonmetals have high ionization energies and tend to form – ions
.
Noble gases: do not form ions
Mostly +2 & +3


Ionization energies increase across the period, because the atoms are getting smaller and the electrons closer to the nucleus and harder to pull away.
Increases


Ionization energies decrease moving down the group because the size of the atom is larger so the electrons are farther away and easier to remove.
In other words the outer electrons are shielded from the pull of the nucleus by the inner shells decreases

E.A. the amount of energy released when an atom gains an e .
The opposite process of ionization. (removing of e )
“affinity” means fondness.
~ E.A. increases across a period. Why? Nonmetals give off more energy when they gain an e than metals do.
~ E.A. decreases down a family. Why? b/c of the larger # of e less energy is given off w/the addition of one more e .


Indicates the elements relative ability to gain electrons in a chemical bond. The greater the electronegativity the greater the attraction for electrons.

How strongly the element “wants” the electron, so metals are low; nonmetals are high


Fluorine is the most electronegative element and Francium is the least. Increases across the periods and decreases down the group. increases
F
0 decreases
Fr


Two most important trends are atomic radius and electronegativity.


In summary:
*Atomic number: a period a family
*Atomic radius: a period
*Ionic radius:
Positive cations
Negative anions a family vs. neutral atom vs. neutral atom

*Ionization Energy: a period
*Electron Affinity: a period
*Electronegativity: a period a family a family a family

Some specific groups:
Group IA Alkali metals : Li, Na, K, Rb, Cs, Fr

In pure state have a silvery appearance and are soft enough to cut with a knife. Yet, alkali metals are so reactive they are not found in nature as free elements. Combine vigorously w/most nonmetals. Usually stored in Kerosene.

Group IIA Alkaline earth metals : Be, Mg, Ca,
Sr, Ba, Ra

Have 2e in outermost level. They are harder, denser, and stronger than alkali metals.
Although they are less reactive than group IA, they are too reactive to be found alone in nature.

Group VII A Halogens : F, Cl, Br, I, At

Most reactive nonmetals. React vigorously w/most metals to form a type of compound known as salts. Has 7e in its outer energy level.

Group VIII A Noble Gases : He, Ne, Ar, Kr,
Xe, Rn

These are the least reactive of all the elements.
Have very stable electron configurations. For many years, the noble gases were believed to be chemically unreactive, yet in the lab inert gas compounds can be synthesized.

1.
On the Periodic table above label:
1) Alkali Metals, 2) Alkaline Earth Metals,
3) Halogens, 4) Noble Gases, 5) Metals,
6) Nonmetals, 7) Metalloids, 8) Transition Metals,
9) Inner Transition Metals

Inner Transition Metals
(f – block elements)
Noble
Gas
Metalloids
(on the stair-step line)

2. Explain why the word periodic is applied to the table of elements.
The pattern of properties repeats periodically with each new row

3. Why do elements in a group (vertical column) in the periodic table exhibit similar chemical properties?
They have the same arrangement of valence electrons

4. What chemical property is common to the elements in group 8A. Explain why.
Group 8A = noble gases = chemically unreactive because they have a full set of 8 (octet) valence electrons

5. In terms of electron configuration, what does the group number of the A-groups tell you?
The A-group number = the number of valence electrons

6. Describe the relationship between the electronegativity value of an element and the tendency of that element to gain or lose electrons when forming a chemical bond.
The higher the electronegativiy, the greater the tendency to gain an electron.

7. Describe the group and period trends in the following atomic properties
Decrease
Increase

7. Describe the group and period trends in the following atomic properties
Increase
Decrease