THE PERIODIC LAWCHAPTER 5

History of the Periodic Table

• In September 1860, scientists gathered together for the First International Congress of Chemists to settle the issue of atomic mass.

• Italian scientist Stanislao Cannizzaro determines an efficient way to measure atomic mass without controversy.

Mendeleev & Chemical Periodicity

• Russian chemist Dmitri Mendeleev accepts atomic mass values discussed at the First International Congress of Chemists and plans on including them in a new chemistry textbook written by him.

• Mendeleev begins to arrange the elements by their properties. However, he notices that when the elements are placed in order of increasing mass, the properties appeared in regular intervals.

• These properties that appear in regular intervals are called periodic.

Mendeleev’s Periodic Table

• Mendeleev created a periodic table in which elements with similar properties were grouped together.

• There were several blank spaces left in Mendeleev’s periodic table. Mendeleev predicted elements would be discovered that would be placed in those spaces.

Moseley and the Periodic Law

• English scientist Henry Moseley discovers a new pattern regarding elements of the periodic table. Moseley finds that the elements fit better when arranged according to increasing nuclear charge (the number of protons in the nucleus).

• This discovery is credited with the development of the term atomic number.

• Moseley’s work was in line with Mendeleev’s arrangement according to properties. Moseley develops periodic law.

Moseley and the Periodic Law

• According to Moseley, the physical and chemical properties of the elements are periodic functions of their atomic numbers.

• Basically, the regularly repeating patterns are due to their atomic numbers.

The Modern Periodic Table• The periodic table of today is an

arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column.

• The periodic table is composed of the s block, p block, d block, and f block. These blocks are based upon the atomic orbitals (or sublevels) that their electrons are filling.

The s-Block Elements• Elements within the s-block are

generally reactive metals.

• Group 1 metals (also called alkali metals) are the most reactive elements. Typically, alkali metals are stored in kerosene to prevent reactions with air or moisture.

• Group 2 metals (also called alkaline-earth metals) are harder, denser, and stronger than alkali metals.

The d-Block Elements

• Elements within the d-block are metals with typical metallic properties.

• These elements are also called transition elements and have a high luster and high conductivity.

• Members of this block are located in Groups 3-12.

The p-Block Elements• The p-block elements consist of all the

elements of Groups 13-18 except helium.

• All members of the p-block have 2 electrons in the s sublevel.

• Elements of the s-block and p-block together are called the main-group elements.

• These elements of the p-block consists of nonmetals, metalloids, and a few metals.

• The halogens (members of Group 17) are the most reactive nonmetals.

The f-Block Elements

• Members of the f-block are located between Groups 3 and 4 in the sixth and seventh periods.

• Lanthanides (elements 58-71) are shiny metals similar in reactivity to the alkaline-earth metals.

• Actinides (elements 90-103) are all radioactive. The first four are found naturally on Earth; whereas, the remaining actinides are synthetically made.

Periodic Properties of the Periodic Table

• There are 5 major periodic properties found on the periodic table:

• Atomic radii

• Ionization energy

• Electron affinity

• Ionic radii

• Electronegativity

Atomic Radii

• Atomic radius = 1/2 the distance between the 2 nuclei of identical bonded atoms

• As you go down a group of elements, AR increases.

• As you go across a period of elements, AR decreases.

Ionization Energy• Ionization energy = the energy needed

to remove an electron from a neutral atom (also known as the 1st ionization energy)

• IE is a measure of how easy an element can become a cation (+ charged ion).

• As you go down a group of elements, IE decreases.

• As you go across a period of elements, IE increases.

Electron Affinity

• Electron affinity = the energy change that occurs when a neutral atom receives an electron

• EA is a measure of how easy an element can become an anion (- charged ion).

• As you go down a group, EA is slightly negative.

• As you go across a period, EA is very negative.

Ionic Radii• Follow same trend patterns as atomic radii!

Electronegativity• Created by American scientist Linus

Pauling

• Scale: 0.0-4.0 (4.0 = fluorine)

• Electronegativity = the ability an element to attract electrons from the electron’s point of view

• As you go down a group, EN decreases or stays the same.

• As you go across a period, EN increases.