Chapter 5 The Periodic Law

Section 5-1 History of the Periodic Table

Stanislao Cannizzaro 13 Jul 1826 – 10 May 1910

Cannizzaro Was the first scientist to accurately measure atomic masses. This was incredibly important for the work that Dmitri Mendeleev was going to do.

Dmitri Mendeleev 8 Feb 1834 – 2 Feb 1907

Dmitri Mendeleev Arranged the elements on cards in order of increasing atomic mass Found columns of elements with similar properties There were gaps in his columns

Basic Version of Mendeleev’s Periodic Table

The gaps Hypothesized the gaps were undiscovered elements. Predicted the props of these elements. Predicted them well.

Mendeleev’s Mistake There were irregularities when arranged according to atomic weight.

Henry Mosely 23 Nov 1887 – 10 Aug 1915

Henry Mosely Discovered a unique charge on the nucleus of the atom. Arranged the elements according to increasing atomic number When he did this, the irregularities disappeared.

Periodic Law The properties of elements tend to change with atomic number gradually, in a periodic way.

John Strutt 12 Nov 1842 – 30 Jun 1919 William Ramsay 2 Oct 1852 – 23 Jul 1916

Strutt & Ramsay In 1894, they discovered argon. Nobody noticed it before because it is completely unreactive. In 1868, helium had been discovered as part of the sun and in 1895, Ramsay showed its existence on earth. In 1898, Ramsay discovered krypton and xenon. Friedrich Ernst Dorn discovered radon in 1900.

Lanthanides Discovered in the early 1900’s. They are found in the f block They are shiny and act like the alkaline earth metals.

Actinides All of these elements are radioactive. They are found in the f block

Alkali Metals Group 1 These elements are soft and can be cut with a knife. They are highly reactive. The will react with both air and water. They form alkaline/basic solutions (the opposite of acidic solutions). Their electron configurations all end s 1.

Sodium

Alkaline Earth Metals Group 2 on the periodic table. These elements are harder and denser than the alkali metals. They are also reactive, but less so than the alkali metals. They will also form alkaline/basic solutions. Their electron configurations all end s 2.

Magnesium

Hydrogen This element doesn’t belong with any group. Its electron configuration is 1s 1.

Helium Even though its electron configuration ends s 2, it isn’t an alkaline earth metal. It is a noble gas because its highest energy level orbitals are full.

Transition metals AKA transition elements AKA d block elements These elements are what we typically picture as common metals. Their d orbitals are being filled.

Transition metals cont’d They are shiny and good conductors of electricity. They are less reactive than the other metals. Some like gold are highly unreactive.

Chromium metal

Main Group Elements Properties of these elements vary greatly because they include metals, nonmetals, metalloids, and noble gases. They include the elements of the s and p blocks. Sometimes they are called the representative elements because metals, nonmetals, metalloids, and noble gases are all represented

Halogens These are the most reactive nonmetals. They have 7 electrons in the outermost energy level and their electron configurations all end in s 2 p 5. They will react with metals to form salts.

Noble Gases Group 18 No stable compounds for He, Ne, or Ar Very low reactivity for the rest Full s & p orbitals (s 2 p 6 ) in the higest energy level –This is very stable - they have no need to react with anything else.

Noble Gases cont’d Most other atoms gain/lose e - to achieve this e - configuration Ne & Ar are used in signs He - Low density - Air ships & weather balloons

Noble Gases Cont’d Rn - Radioactive Found in homes - Linked to Lung CA Once you test for it, you must disclose the results to potential buyers

5 – 3 Electron Configuration and Periodic Properties

Atomic Radius As you move down a group it increases The outermost e - are being added to higher energy levels (further from the nucleus.

Atomic Radius As you move across a period, it decreases Even though e - are being added, they are added to the same energy level (same distance from the nucleus).

Atomic Radius The charge on the nucleus increases as you move across the period and so it has a “tighter” hold on the e - being added.

Shielding Effect The reduction of the attractive force between a nucleus and its outer electrons due to the blocking effect of inner electrons.

Ionization Energy The amount of energy needed to remove an electron from an atom

Ion An atom that has gained or lost an e - If it has gained an e -, it will be _____. If it has lost an e -, it will be _____.

Ionization Energy As you go down a group, it decreases Shielding effect and electrons are being added to higher energy levels.

Ionization Energy As you move across a period, it increases The charge on the nucleus increases as you move across the period and so it has a “tighter” hold on the e - being added.

Electron Affinity Measures the tendency of an atom to attract electron The energy change that occurs when an electron is acquired by a neutral atom. Metals tend to have positive energy changes, they do not have a tendency to attract electrons Non-metals tend to have negative energy changes, they have a strong tendency to attract electrons

EA As you move down a group, electrons add with greater difficulty (values become more positive) Shielding Effect and electrons are being added to higher energy levels. There are exceptions

EA Electrons add more easily as you move across the periodic table.(values become more neg) The charge on the nucleus increases as you move across the period and so it has a “tighter” hold on the e - being added.

Electronegativity Tendency for an atom to attract e - to itself when combined with another atom. F is the most EN EN decreases as you move down a group EN increases as you move across a period

Electronegativity Based on the Pauling Scale. Linus Pauling 28 Feb 1901 – 19 Aug 1994