The Periodic Law

History of the Periodic Table

Problems

• Different chemists used different masses for elements.

• There were no accurate ways to determine the number of atoms in a compound.

• Due to different masses, there were different compositions calculated.

Mendeleev and Chemical Periodicity• Used values of atomic masses

established by Cannizzaro.• Wished to organize elements by their

properties. • Looked for trends or patterns in

properties. • Found repetitive properties---

periodicity---when elements were arranged according to mass.

Mendeleev and Chemical Periodicity• Grouped elements with similar

properties together. • Created the first periodic table.• Left open spaces for predicted

elements.• Note: The elements predicted by

Mendeleev in 1871 were found by 1886 and their properties were almost exactly those predicted by Mendeleev years earlier.

Properties of Some Elements Predicted by Mendeleev

Moseley and the Periodic Law• Organized elements by increasing

nuclear charge. • Led to the atomic number concept.• Reorganized the periodic table in order

of atomic number.• Supported Mendeleev’s periodic law

which stated “The physical and chemical properties of the elements are periodic functions of their atomic numbers.”

The Modern Periodic Table• It is an arrangement of the elements in

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

• The most significant additions were the noble gases. Helium was discovered in 1868. Argon was discovered by Ramsey in 1894. He also discovered krypton and xenon in 1898. Dorn discovered radon in 1900.

The Modern Periodic Table

• There were additions to the table after Mendeleev’s time.

• The noble gases---They are known by this name because they are nonreactive or inert.

• The lanthanides---elements 58 to 71.• The actinides---elements 90 to 103.

The Periodic Law

Electron Configurations and the Periodic Table

Periods and Blocks of the Periodic Table

• In the first period, the 1s sublevel is filled. Since it holds only 2 electrons, the first period must fill with 2 electrons.

• In the second period, there are ___ sublevels, ___ and ___.

• In the third main level, there are ___ sublevels.

Periods and Blocks of the Periodic Table• The period of an element can be found

by its configuration. Arsenic, for example, is [Ar]4s23d104p3 which means that it is in the fourth period because the highest occupied main level is n = 4.

• Based on the electron configurations of the elements, the periodic table can be divided into four blocks, the s, p, d, and f blocks.

Blocks of the Periodic Table

The s-Block Elements: Groups 1 and 2• These are chemically reactive metals. The

Group 1 (ns1) metals are more reactive than the Group 2 (ns2) metals.

• Group 1 elements contain ___ valence electron(s) and are known as the alkali metals.

• Group 2 elements contain ___ valence electron(s) and are known as the alkaline earth metals.

• Hydrogen and Helium are in the s-block but their properties are not like the others in the block. They are unique.

The d-Block Elements: Groups 3-12The Transition Metals• The d-sublevel first appears in the ___ main

level. • The d-sublevels are slightly higher in energy

than the previous s-sublevel. This is where the main energy levels begin to overlap.

• There are exceptions in this block. Some elements fill electrons to achieve the lowest possible energy state. Two exceptions in the 4th period are chromium and copper. Experimentally it was found that sublevels prefer to be half-filled or completely filled if they are close to either state.

The d-Block Elements: Groups 3-12• Chromium has the configuration

1s22s22p63s23p64s13d5. Molybdenum would be similar.

• Copper has the configuration 1s22s22p63s23p64s13d10. Silver would be similar.

• Palladium, platinum, and many others are also exceptions to the normal filling pattern. In each case the abnormal filling occurs to provide the lowest possible energy state for the atom.

• These elements are not as predictable as the s-block or the p-block. They usually have more than one oxidation state.

Exceptional Electron Configurations• Niobium [Kr] 5s14d4

• Molybdenum [Kr] 5s14d5

• Technetium [Kr] 5s14d6

• Ruthenium [Kr] 5s14d7

• Rhodium [Kr] 5s14d8

• Palladium [Kr] 4d10

• Silver [Kr] 5s14d10

• Platinum [Xe] 6s1 4f14 5d9

The p-Block Elements: Groups 13-18• These are atoms which add an electron to the

p-sublevel. • These are main-group elements or

representative elements. • All of the nonmetals and metalloids are

included in the p-block. There are 8 metals here also.

• The halogens(Gp 17) are the most reactive nonmetals. Fluorine is the most reactive nonmetal on the periodic table.

The f-Block Elements: Lanthanides and Actinides• These are located on the bottom of the

periodic table but are really placed in the sixth and seventh periods.

• The first main level that has an f-sublevel is the ___ level.

• There are 14 elements on each row. • They are similar in reactivity to Group 2

elements.

Main Group (Representative) Elements• Group 1 – Alkali metals – very reactive metals

– never found freely in nature• Group 2 – Alkaline earth metals – reactive

metals – never found freely in nature• Group 3(13) – Boron group – less reactive –

contains a metalloid and metals• Group 4(14) – Carbon group – most diverse

group• Group 5(15) – Pnictogens• Group 6(16) – Chalcogens• Group 7(17) – Halogens – most reactive

nonmetals • Group 8(18) – Noble gases – most stable

Relationship between Periodicity and Electron Configuration

Electron Configuration and Periodic Properties

Trends Exhibited on the Periodic Table

Atomic and Ionic Radii• The atomic radii of elements decrease

as one moves across a period. (Why?)• This is caused by an increasing nuclear

positive charge while electrons are added into the same main level.

• The atomic radii of elements increase as one moves down a group.(Why?)

• This occurs because more energy levels are being filled as one goes down a group.

Trends in Atomic Radii

Atomic Radii

• Going down groups, the radii of transition elements increase from the 4th to the 5th

period. However, the 6th period transition elements, are about the same size as the 5th period transition elements.

• This is due to the lanthanide contraction, the placement of the lanthanides after the 6s sublevel. The expected increase in size does not occur.

Trends in Atomic Radii

Ionic Radii

• Atoms form ions by losing or gaining electrons.

• An atom that has lost an electron becomes ___ and an atom that has gained an electron becomes ___.

• A positive ion is called a ___.• A negative ion is called an ___.

Isoelectronic Character• Isoelectronic means that species will have the

same number and distribution of electrons. Ions will have the same electron distribution as the nearest noble gas.

• There is an isoelectronic series on every row. • The more positive ion in such a series is the

smallest. The more negative ion in such a series is the largest.

Isoelectronic Series• An isoelectronic series is Al+3, Mg+2,Na+1, Ne, F-1,

O-2, and N-3. All are 1s22s22p6. Which of the above ions is largest?

• N-3 is. Why is that so? • Nitrogen has 7 protons and 7 electrons, but

when the ion is formed, 3 more electrons are added. There is an imbalance in the positive and negative charges. There are more negatives than positives so the electron cloud is not held as tightly. In this species there are the fewest protons for the 10 electrons, so it is the largest.

Ionic Radii

Trends in Ionization Energies

Ionization Energy

• Ionization energy is the amount of energy required to remove an electron from a gaseous atom. This is always an endothermic change. – A + first ionization energy A+ + e-

• It is usually expressed in kJ/mol.• The value of the first ionization energy

depends on the effective nuclear charge, the atomic radius, and the electron configuration.

Ionization Energy• I.E. increases with each successive electron

removal. (Why?)• As each electron is removed there is more

nuclear charge on each electron remaining with the atom. Therefore more energy is required.

• The amount jumps dramatically when trying to remove an electron from a filled, stable level.

Ionization Energy• Which of the following species would

require the most energy to remove a first electron? Rb or I

• Iodine because there are more electrons between the outer electron and the nucleus but the single electron of Rb does not have that protection(inner electron shielding).

• I.E. increases across to the Noble Gases and decreases down a group.

Electron Affinity• Neutral atoms can also acquire electrons. • The energy change that occurs when an

electron is acquired by a neutral atom is called the electron affinity.

• A + e- A- + energy or A + e- + energy A- may be the general equation. The change may be endothermic or exothermic, but is exothermic most of the time.

Electron Affinity• Elements in the upper right corner (Gp 17)

tend to have large negative electron affinities meaning that they become more stable by gaining an electron. That would be a high electron affinity.

• Elements in the lower left corner tend to have much less negative electron affinities. They are unlikely to gain an electron.

Electron Affinity

Electronegativity

Electronegativity

• Electronegativity is a measure of an atom’s ability to attract electrons in a compound.

• Fluorine is the most electronegative element on the periodic table.

• Electronegativity decreases as one moves down and to the left.

Summary of Trends