The Periodic Table and The Periodic Table and PeriodicityPeriodicity

• Mendeleev’s brilliant organizational skills• The modern table – groups, families and series and

trends

Areas of InterestAreas of Interest

The father of the periodic table.

In the 19th century elements were being discovered rapidly, a way was need to organize them.

He arranged the atoms according to increasing atomic weight.

Ok so what?

The brilliance of his arrangement came from the atoms he left off the table . . . Those that had not yet been discovered.

Mendeleev arranged his table in rows and columns that not only addressed increasing atomic mass but was able to predict undiscovered elements based on properties.

Dimitri MendeleevDimitri Mendeleev

For years chemists had known about elements sharing similar properties, in 1869 Dimitri Mendeleev provided an organized arrangement.

His most famous omission he named eka-silicon. He predicted an element that had a greater mass than silicon, a smaller mass than tin but shared similar properties with both elements.Property Ekasilicon Germanium

atomic mass 72 72.59

density (g/cm³) 5.5 5.35

melting point (°C) high 947

color gray gray

oxide type refractory dioxide refractory dioxide

oxide density (g/cm³) 4.7 4.7

oxide activity feebly basic feebly basic

chloride boiling point under 100°C 86°C (GeCl4)

chloride density (g/cm³) 1.9 1.9

Henry Moseley determined the atomic number of each element and placed them in order of increasing atomic number, rather than atomic weight.

This atomic number arrangement also followed Mendeleev’s trends in properties. Elements with similar properties were grouped together.

The modern period table was born (1914).

Arranged in rows and columns.

A row is called a period

A column is called a group or family

Some of the groups (or families) have special names

that help us identify them as a collective.

Famous families if you will . . .

The include all of Group 1: Li, Na, K, Rb, Cs and Fr*.

Soft shiny metals that react violently with water to produce H2 gas.

Electron configurations of ns1. (ie Li is 1s22s1)

Readily form +1 cations (ie Na loses and electron to form Na+)

*Francium only exists for microseconds so it cannot be studied in quantity.

The include all of Group 2: Be, Mg, Ca, Sr and Ba.

These metals are soft but not quite as much as those of Group 1.

They are stable in air (unlike the Alkalai Metals)

Electron configurations of ns2. (ie Be is 1s22s2)

Readily form +2 cations (ie Mg loses 2 electrons to form Mg2+)

Electron-rich elements that most resemble what we think of when we talk about metals:

they’re malleable and ductile

they conduct electricity

the free flow of electrons yields many colorful solutions

they’re shiny

they conduct heat

Many of these elements are synthetic, they’re made in particle accelerators and used for research or highly specific purposes.

They are metals but they are very dense and many are quite rare.

Elements include: B, Si, Ge, As, Sb, Te and At

They’re not quite metals but they’re not quite non-metals.

They’re semi-conductors (they can selectively conduct electricity).

The metalloids

Si, the semiconductor the computer industry is built upon

The Halogens

Elements include: F, Cl, Br, I and At

Readily form -1 anions (ie Cl gains an electron to form Cl-)

React well with metals from Groups 1 and 2.

Behave as other non-metals (non-conductive, not shiny etc.)

These are the elements found in Group 18, the farthest to the right on the periodic table.

They are all gases and are VERY stable (they do not readily undergo reaction).

The have full energy levels and sub-shells. For example Ar has electron configuration 1s22s22p63s23p6.

When we pass a high-voltage current through any of these gases we get extremely bright light.

Metals Transition MetalsLanthanides & ActinidesMetalloids and Non Metals

Use your textbook to define the following termsUse your textbook to define the following terms

1. Atomic radius – half the distance between 2 nuclei in a covalent bond (an estimate of the radius of a single atom)

2. 1st Ionization Energy – the energy required to remove the 1st electron from an atom in the gaseous state (generating a +1 cation)

3. 2nd Ionization Energy – the energy required to remove the 2nd electron from a gaseous atom (generating a +2 cation)

4. Electron Affinity – the energy released when a gaseous atom captures an electron (generating a -1 anion)

5. Electronegativity – the tendency of an atom to attract electrons to itself (higher electronegativity means the atom “wants” electrons)

6. Activity – a trend that dictates replace-ability in single replacement reactions (decreases down the halogens)

7. Valence Electron – an electron in the highest occupied energy level of an atom (these are the ones involved in bonding)

8. Effective Nuclear Charge – the “pull” on the electrons from the protons in the nucleus (decreases as atoms get larger)

Why this pattern?

Decreasing effective nuclear charge (the “pull” the electrons feel from the protons in the nucleus)

Atomic Radius (Atomic Size)

Atomic Radius – Half the distance of a pure covalent bond

Atomic Radius (Atomic Size)

1st Ionization Energy

- the energy required to remove 1 electron from a neutral atom

Ionization energy decreases down a Group.

Ionization increase from left to right in a period.

Increasing Electronegativity

Increasing Effective Nuclear Charge

Your AssignmentYour Assignment

1. This element is larger than chlorine but smaller than iodine.2. Of the following pairs, which is MORE metallic:

a. Si or Geb. As of Gec. Ba or Csd. Be or Be. Kr or Xe

3. The electron configurations of six neutral atoms are shown:i. 1s2 2s2 2p6 3s1

ii. 1s2

iii. 1s2 2s2 2p6 3s2

iv. 1s2 2s2 2p6

v. 1s2 2s1

vi. 1s2 2s2 2p3

a. Which has the highest first ionization energy?b. Which has the lowest first ionization energyc. Which would have the lowest second ionization energy?d. Which would likely have the lowest third ionization energy?