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.