Chapter 6

The Periodic Law

Objectives: Describe the periodic tables of Moseley and

Mendeleev.

Identify the various families of elements on the periodic table.

State the trends in atomic radius, ionization energy, electron affinity and ion size with a group or period on the periodic table.

Identify the relationship between these trends and the structure of the atom.

Chapter 6Section 1

History of the Periodic Table

By 1860, more than 60 elements had been discovered.

Chemists needed to learn the properties of these elements as well as those of the compounds they formed.

At the time there was no method for accurately determining an element’s atomic mass.

Things were disorganized and confusing among chemists.

Dmitri Mendeleev

Mendeleev hoped to organize the elements according to their properties.

He arranged the elements in order of increasing atomic mass and noticed certain similarities in their chemical properties would appear at regular intervals.

In 1869 Mendeleev created the first periodic table grouping the elements with similar properties.

Henry Moseley

In 1911 Moseley revised the periodic table of Mendeleev by arranging the elements in order of increasing atomic number (number of protons) and not atomic mass.

The elements were still grouped by similar physical and chemical properties.

Modern Periodic Table

The periodic table has undergone extensive changes since Mendeleev’s time.

More than 40 new elements have been discovered or synthesized.

Periodic Table – an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group.

Periodic Table

Elements are classified in three major groups: Metals Nonmetals Metalloids

Use the periodic table to distinguish the classes of elements

Physical Properties of Metals

Malleability Ductility Luster Heat conductors and

electrical conductors Solids Ex. Iron (Fe), tin (Sn),

zinc (Zn), and copper (Cu)

Properties of Nonmetals

Dull in appearance Brittle Do not conduct electricity Ex. Carbon (C), Oxygen (O)

and Sulfur (S)Solids, liquids or gases

The metalloids divide the metals from the nonmetals.

They are mostly brittle solids with some properties of metals and some of nonmetals.

The electrical conductivity falls between the metals and nonmetals.

The metals of the p block are reactive enough to be found in nature only as compounds and not free elements (except bismuth).

Metalloids

Properties of metals and nonmetals

Ex. Silicon (Si) and Germanium (Ge)

Common in the computer

industry

Summary

Recognize the work of Mendeleev and Moseley.

How the modern periodic table is arranged with respect to the elements.

Know how the periodic table is divided into different groups.

Homework

Page 198

Questions 29, 30 , 31 and 32

Chapter 6Section 2

Electron Configuration and the

Periodic Table

The elements arranged vertically in the periodic table share chemical properties due to the same number of valence electrons.

They are also organized in horizontal rows, or periods.

The length of each period is determined by the number of electrons that can occupy the sublevels in that period.

The period of an element can be determined from the element’s electron configuration.

For example, arsenic – As, has the electron configuration of:

[Ar]3d104s24p3

The 4 in 4p3 indicates that arsenic is in the fourth period of the periodic table.

The periodic table is divided into 4 blocks

Without looking at the periodic table give the period number for the following electron configuration:

[Xe]6s2

The s-Block ElementsThe elements of the s block are chemically reactive metals.

The elements of the s-block and p-block are known as the main-group elements.

The outermost energy level in an atom of the Group 1 elements contain a single electron.

The ease with which the single electron is lost make the Group 1 metals extremely reactive.

Alkali Metals – The elements in Group 1 of the periodic table (lithium, sodium, potassium, rubidium, cesium and francium).

Because the alkali metals are so reactive they are not found in nature as free elements.

They combine vigorously with most nonmetals.

They react strongly with water to produce hydrogen gas.

Alkaline-Earth Metals – The elements in Group 2 of the periodic table (beryllium, magnesium, calcium, strontium, barium and radium).

Atoms of the alkaline-earth metals contain a pair of electrons in their outermost shell.

Slightly less reactive than the alkali metals.

Still too reactive to be found in nature as free elements.

The d-Block ElementsThe elements of the d block are known as the transition metals.

They are good conductors of electricity and have a high luster.

They are typically less reactive than the alkali and alkaline-earth metals.

Palladium, platinum and gold are among the least reactive of all elements.

The p-Block ElementsThe elements of the p-block and s-block are known as the main-group elements.

The p-block consists of the elements in Groups 13-18 (except helium).

The properties of the p-block vary greatly.

The p-block contains metals (Al), metalloids (Si) and nonmetals (Br).

The electron configuration of bromine is:

[Ar]3d104s24p5

Halogens – the elements of Group 17 (fluorine, chlorine, bromine, iodine and astatine).

The halogens are the most reactive nonmetals.

They react vigorously with most metals to form compounds known as salts (NaCl).

Fluorine and chlorine are gases. Bromine is a liquid and iodine is a solid.

Noble Gases (1868-1900) – group 18 elements that are characterized by their relative un-reactivity.

The Lanthanides (early 1900’s) – the 14 elements with atomic numbers from 58 (cerium, Ce) to 71 (lutetium, Lu)

The Actinides – the 14 elements with atomic numbers from 90 (thorium, Th) to 103 (lawrencium, Lr).

The lanthanides and actinides belong in periods 6 and 7 of the periodic table.

To save space and to group them together they are set off below the main portion of the periodic table.

The Group 18 elements (Noble Gases) undergo few chemical reactions.

This stability results from the gases electron configuration.

Their highest occupied levels are completely filled with electrons (octet).

An atom’s electron configuration governs the atom’s chemical properties.

Classwork

Page 186

Problems: 12, 13 and 14

Homework

Page 198

Questions 43, 44, 47 and 48

Chapter 6Section 3

Electron Configuration and the

Periodic Properties

So far you have learned that the elements are arranged in the periodic table according to their atomic number.

There is also a correlation between the arrangement of the elements and their electronic configuration.

We will look at the relationship between the electron configurations and the periodic trends of the elements.

Ideally, the size of an atom is defined by the edge of its orbital.

However, this boundary varies under different conditions.

One way to express the atomic radius is to measure the distance between the nuclei of two identical atoms that are bonded together, then divide this distance by two.

Atomic Radii

Atomic radius may be defined as one-half the distance between the nuclei of identical atoms that are bonded together.

Period Trends

There is a gradual decrease across a row.

There is a gradual decrease across a row.

The trend to smaller atoms across a period is caused by increasing positive charge of the nucleus. Adding of protons or increasing atomic number.

The electrons are pulled closer to the nucleus.

The increase pull results in a smaller atomic radius.

Group Trends

There is an increase down a group.

There is an increase down a group.

As electrons are added to sublevels in higher energy levels located further from the nucleus, the size of the atoms increase.

An exception is between aluminum (radius – 143 pm) and gallium (radius – 135 pm).

This is due to the gallium being proceeded by the 10 d-block elements.

The nuclear charge is considerably higher.

Problem

Of the elements magnesium-Mg, chlorine-Cl, sodium-Na, and phosphorus-P, which has the largest atomic radius and why?

Of the elements calcium-Ca, beryllium-Be, barium-Ba and strontium-Sr, which as the largest atomic radius and why?

Classwork

Page 189

Problems: 16 and 17

Ionization energy (IE) – the energy required to remove one electron from a neutral atom of an element.

A + energy A+ + e-

An ion is an atom or group of bonded atoms that has a positive or negative charge.

Ionization Energy

Period Trends

In general, ionization energies of the elements increase across a period.

Group 1 elements – have the lowest ionization energies. Therefore they lose electrons most easily. Very reactive.

Group 18 elements - have the highest ionization energies . They do not lose electrons easily. Very low reactivity.

The increase is due to increasing nuclear charge (more protons going across a period).

A higher positive charge more strongly attracts electrons in the same energy level.

Therefore, it is tougher to remove an electron from an atom.

Increasing nuclear charge is responsible for both an increasing ionization energy and decreasing atomic radius across a period.

Group Trends

Ionization energies generally decrease down a group.

Electrons removed from atoms of the elements down a group are farther from the nucleus.

Also, the electrons from the lower energy levels shield the outer electrons.

Therefore, they are removed more easily.

Neutral atoms can also acquire electrons.

Electron Affinity – the energy change that occurs when an electron is acquired by a neutral atom.

A + e- A- + energy

Electron Affinity

Period Trends

Among the elements of each period, the halogens (Group 17) gain electrons most readily.

The ease with which halogen atoms gain electrons is a major reason for the high reactivities of the Group 17 elements.

As electrons add to the same p sublevel of atoms, electrons affinities increase.

Period Trends

An exception to this trend occurs between Group 14 and 15.

Compare the electron affinities of carbon ([He] 2s2 2p2 ; -126.3) and nitrogen [He] 2s2 2p3; 0).

Adding an electron to nitrogen is more difficult because it forces the electron to pair with another electron.

Group Trends

Trends for electron affinities within groups are not as regular as trends for ionization energies and atomic radii.

Cation (A+) – a positive ion that results with loss of an electron. Related to ionization energy.

Cations and Anions

Anion (A-) – a negative ion that results with the gain of an electron. Related to electron affinity.

Valence Electrons

Chemical compounds form because electrons are lost, gained or shared between atoms.

The electrons that interact are those in the highest energy levels.

Valence electrons – the electrons available to be lost, gained or shared in the formation of chemical compounds

For the main group elements, the valence electrons are the electrons in the outermost s and p sublevels.

The inner electrons are in filled energy levels and are held to tightly to be involved.

Electronegativity

Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons.

Period Trends

Electronegativities tend to increase across a period.

The Group 1 and 2 elements have the lowest electronegativities.

Groups 16 and 17 have the highest electronegativities.

Group Trends

Electronegativities tend to decrease down a group.

The combination of these trends in electronegativities results in the highest values belonging to the elements in the upper right of the periodic table (fluorine).

Problem

Of the elements gallium-Ga, bromine-Br, and calcium-Ca, which has the highest electronegativities and why?

Homework

Page 199

Questions 58, 60, 63 and 64

Due: