Chapter 5

The Periodic Law

I: History of the Periodic Table

by 1860, sixty elements had been discovered

1860, a gathering of chemists in Germany


A: Mendeleev and Chemical Periodicity

use the new atomic masses and organize elements according to their properties

similarities in their chemical properties appeared at regular intervals

periodic - definition

Figure 5-1 page 123

elements with similar properties were grouped together

atomic masses as a guide but he let the properties determine the exact grouping of elements e.g. Tellurium (Te) and Iodine (I)

empty spaces for elements

Two questions remained

B: Moseley and the Periodic Law

Henry Moseley working with Ernest Rutherford

arranged in increasing order according to nuclear charge

atomic number

explained the tellurium/iodine question

Periodic Law - statement

C: The Modern Periodic Table

Periodic Table - definition

1. The Noble Gases

1894 - John Strutt and William Ramsay

1868 helium; 1895 - Ramsay

group 18

1898 - Ramsay

1900 - Friedrich Dorn

2. The Lanthanides

lanthanides - definition

similar in chemical and physical properties

part of period 6

3. The Actinides

part of period 7

at the bottom to save space

normally fit between groups 3 and 4

4. Periodicity

Figure 5-4 page 126

The differences in atomic numbers for the two groups follow the same pattern

key is looking for the patterns

Homework: 5.1

II: Electron Configuration and the Periodic Table

Group 18 - octet - stable

Configuration of valence electrons determines the atom's chemical properties

A: Period and Blocks on the Periodic Table

Vertical columns (group or family) have similar chemical properties

Length of each period is determined by the number of electrons that can occupy the sublevels being filled in that period - determined by the electron configuration of the elements

Table 5-1 page 128

1st period 2 elements filling s sublevel

2nd period: 8 elements filling s and p sublevel

3rd period: 8 elements filling s and p sublevel

4th period: 18 elements filling s and d and p sublevels

Figure 5-5 page 129

1. The s Block Elements: Groups 1 and 2

active metals - group 1 being more active than group 2

Group 1 = alkali metals

Group 1: each element has one electron in an s sublevel

Tend to lose that electron to form a +1 ion and achieve an octet

In the elemental state (uncombined) they have a silvery appearance and can be cut with a knife

Not found in nature in the elemental (uncombined) state. Instead they react vigorously with nonmetals and form compounds - the way they are found in nature

React strongly with water to form hydrogen gas and an aqueous solution of alkalis (solutions of a base)

Stored under kerosene to protect them from reacting with the moisture in the air

Moving down the column the melting points decrease

Group 2 = alkaline earth metals

each element has two electrons in an s sublevel

tend to lose two electrons to form a +2 ion and achieve an octet

harder, denser, and stronger than alkali metals

have higher melting points than alkali metals

less reactive than alkali metals

not found in nature in the elemental (uncombined) state

2. Hydrogen and Helium

one electron in the s sublevel

separated from Group 1 metals

Properties of hydrogen do not resemble those of any other group

Helium has two electrons in the s sublevel

part of group 18

ts filled s sublevel gives it special stability

Different than group 2 metals since they have an unfilled p sublevel of the energy level n

3. The d-block Elements: Groups 3-12

After the ns sublevel is filled, electrons go into the (n-1)d sublevel until the d sublevel is filled. After the (n-1)d sublevel is filled the np sublevel is filled.

d-block elements are metals

called transition elements or transition metals

less reactive than alkali metals and alkaline earth metals

some are unreactive and exist in nature in the elemental (uncombined) form e.g. palladium, platinum and gold

4. The p-Block Elements: Groups 13-18

electrons enter the p sublevel only after the s sublevel of the same energy level is filled

main-group elements

properties vary

right side

six metalloids

on the left, bottom

Group 17 - Halogens

most reactive nonmetals

forms salts

fluorine and chlorine




Metalloids or Semiconducting Elements

on red stepped line

mostly brittle solids

metals of p block

generally harder and denser than the s-block alkaline earth metals, but softer and less dense than the d-block metals

except for bismuth, found in nature in combined state

when purified as pure element, they are stable in the presence of air

The f-Block elements - Lanthanides and Actinides

between groups 3 and 4

fourteen f-block elements in each series - how many electrons can an f block hold?

similar in reactivity to the group 2

actinides are all radioactive

the first four of the actinides vs remaining actinides

Homework: Chapter 5: 5.2

III: Electron Configuration and Periodic Properties

periodic law and the electron configurations

A: Atomic Radii

atomic radius - definition

1. Periodic Trends

figure 5-13 page 141

figure 5-14 page 142

across second period

2. Group Trends

group 1

group 13

interplay between the distance of the added electrons from the nucleus and the increased number of protons in the nucleus

B: Ionization Energy

A + I.E. ---> A+ + e-

Ion - definition

Ionization - definition

compare ease with which atoms give up electrons

ionization energy - definition

figure 5-15 page 143

page 144

1. Period Trends

First element in series vs last element - for ionization energy

Low I.E. indicates

High I.E. Indicates

trend across period and why

ionization energy of metals vs nonmetals

2. Group Trends

trend for main group down family and why

3. Removing Electrons from Positive Ions

ions that result from applying first, second and third ionization energy

Size of first vs second vs third ionization energy and why

table 5-3 page 145

ionization energy of group 18 vs other groups

Difference between first and second I.E. of Li

Difference between second and third I.E. Of Be

Difference between I.E. Of beryllium and boron

Difference between I.E. Of nitrogen and that of oxygen

C: Electron Affinity

Electron Affinity - definition

A + e- ---> A- + electron affinity (energy)

Some atoms need to absorb energy to accept the electron

A + e- + electron affinity (energy) --->A-

The ions formed are unstable

figure 5-17 page 147

In tables not equations: energy will have a negative value for exothermic reactions; energy will have a positive value for endothermic reactions

1. Periodic Trends

Halogens gain electron most readily - size of electron affinity?

Why does giving off a large amount of energy indicate the atom wants the electron?

Across the p block of any series, adding an electron produces greater negative values.

Not the case between group 14 and group 15. Why?

2. Group Trends

not as regular as for I.E.

down the group what happens and why?

Many exceptions in the transition metals

3. Adding Electrons to Negative Ions

add an electron to a negative ion - repulsion

all second electron affinities are positive

D: Ionic Radii

Cation - definition

Radius of the ion vs radius of atom and why

Anion - definition

radius of anion vs radius of atom and why

Figure 5-19 page 149

1. Periodic Trends

ions that metals tend to form and nonmetals tend to form

Cationic radii decrease across a period - why

Beginning with group 15, anions tend to form

Anionic radii decrease across the period - why

2. Group Trends

Gradual increase in ionic radius down the group - why

E: Valence Electrons

Valence Electrons - definition

Usually found in incompletely filled main energy levels

Table 5-4 page 150

Elements in groups 13-18 have a number of valence electrons equal to the group number minus 10

F: Electronegativity

valence electrons hold atoms together in chemical compounds

Valence electrons are not always mid-way between the two atoms

This affects the chemical properties of the compound

Electronegativity - definition


relative scale with 4.0 being highest

1. Periodic Trends

figure 5-20 page 151

increase going across the periods - generally

active metals are the least electronegative

active nonmetals are the most electronegative

Down a group the electronegativities tend to decrease down a group or remain about the same - why

Some noble gases do not form compounds and do not have electronegativities assigned - why

Highest electronegativities are in the upper right of the periodic table - why

Figure 5-21 page 152

G: Periodic Properties of the d and f Block Elements

d block elements vary less and with less regularity than the main group elements. e.g. flat curves in figure 5-14 and 5-16

For d block elements, electrons in both the Ns and (n-1)d sublevel are available to interact with the surroundings

Electrons in the incompletely filled d sublevels are responsible for many characteristic properties of the d block elements

1. Atomic Radii

generally decrease across a period

less than that for main group elements

in figure 5-14 the radii decrease then increase slightly across each of the four periods that contain d block elements


2. Ionization Energy

ionization energy generally increase across the periods

ionization energy increases going down group

shielding of s electrons by d electrons

3. Ion Formation and Ionic Radii

Order of removing electrons.

Highest energy level then highest energy sublevel where s < p < d < f

e.g. Fe

Most d block elements form 2+ ions - why

Cations have smaller radii than atoms - why

4. Electronegativity

Values of electronegativity increase as radii decreases and vice versa

Summary of Periodic Trends

Homework: 5.3

end of notes


Periodic means recurring at regular intervals.

Periodic Law state the physical and chemical properties of the elements are periodic functions of their atomic numbers.

The periodic table 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 lanthanides are the fourteen elements with atomic numbers from 58 to 71.

The actinides are the fourteen elements with atomic numbers from 90 to 103.

Ionization energy is the energy required to remove on electron from a neutral atom of an element.

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

Any process that results in the formation of an ion is referred to as ionization.back

Atomic Radius is one-half the distance between the nuclei of identical atoms that are bonded together.

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

A cation is any positive ion.

An anion is any negative ion.

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

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