Chapter 6: The Periodic Table and Periodic Law

Development of the Periodic Table

• 1790s – Antoine Lavoisier composed a list of the 23

known elements– Included gold, silver, carbon, and oxygen

• Electricity which is used to break down compounds into elements led to an “explosion” in chemistry as did the spectrophotometer and the industrial revolution.

• 1864– John Newlands proposed an organization

scheme for the elements– Arranged by increasing atomic mass and

noticed that the properties of the elements repeated after every 8th element (PERIODIC)

– See fig 6.2 page 153 for picture

Meyer, Mendeleev, and Moseley, OH MY!!!

• Lothar Meyer and Dmitri Mendeleev showed a connection between atomic mass and elemental properties

• Mendeleev published first!!!• Left spaces on the Periodic Table for the

unknown elements • By noting trends in the periodic table, he was

able to predict the properties of yet to be known elements.

• Mendeleev organized the periodic table by atomic mass

Mendeleev’s Periodic Table• http://http://z.about.com/d/chemistry/1/0/0/W/mendeleevperiodic.jpg

• Mendeleev was not completely correct – More accurate measurements of atomic mass

• Mosely (1913)- arranged elements in order of increasing atomic number – Resulted in clearer patterns of properties

• PERIODIC LAW:There is a periodic repetition of chemical and physical

properties of the elements when they are arranged by increasing atomic number

The Modern Periodic Table

• Groups - the columns of the periodic table(Sometimes called families)

• Periods – the rows of the periodic table

• SEE PAGE 154 Fig. 6.4

• Representative Elements (labeled 1A-8A)• Transition Elements (labeled 3B-12B)

Classifying the Elements (3 Types)

1. Metals (solid, shiny, good conductors)• Group 1A: Alkali Metals

• Most reactive of all metals

• Group 2A: Alkaline Earth Metals• Also very chemically reactive but not as much as

the alkali metals

• Group 3A: Transition Metals (main part of table) and Inner Transition Metals (bottom two rows, lanthanide and actinide series)

2. Nonmetals – generally gases, dull, brittle, poor conductors– Group 7A is called the halogens and are very reactive – Group 8A is called the Noble Gases and are

unreactive

3. Metalloids – Phys. and chem. properties of both metals and

nonmetals– Border on the stair-step line– Silicon and Germanium are two most important

(comp. chips)

6.2: Classification of the Elements

• Valence electrons • Found in highest principle energy level• All elements of group 1A have the same number of

valence electrons; therefore, have same chemical properties

• Valence electrons by period• The energy level by the valence electrons are

found reveals the period

• Valence electrons by group• The group number corresponds the number of

valence electrons

The s, p, d, and f elements

• Review pages 160-161 in case you had trouble or may be a little confused

6.3 PERIODIC TRENDS

This is a very important section!!!!!

• Many properties on the periodic table change in a very predictable manner

• Includes:– Atomic Radius– Ionic Radius– Ionization Energy– Electronegativity

• YOU MUST MEMORIZE THESE!!!!!!

1. Atomic Radius

• Atomic size is based on how closely an atom is to it’s neighboring atom

• Because the neighboring atom can vary from one substance to another, the size itself tends to vary

• For sodium, The atomic radius is defined as half the distance between adjacent nuclei in a crystal of an element

Atomic Radius CONT’D

• Trends within the periods– As you move left-right, atomic size decreases– Caused by the increasing positive charge in a

nucleus– Each successive element increases in

number of electrons and protons– Remain in same principal energy level– The increased nuclear charge pulls the

outermost electrons in closer to the nucleus

Atomic Radius CONT’D

• Trends within groups– Increase as you move down a group– The nuclear charge increases and electrons are

added to higher principal energy levels– Outer electrons are farther from the nucleus

1

2

3

4 5

6

7

2. Ionic Radius

• Atoms can gain or lose electrons to form ions• Because electrons are negatively charged the

change in quantity causes there to be a change in the net charge

• ION- an atom or bonded group of atoms that has a positive or negative charge– When atoms lose electrons, they become positive

and, therefore, are smaller– When atoms gain electrons, they become negative

and, therefore, are larger

• Ionic Radius– Cations (+)

• lose e-

• smaller

© 2002 Prentice-Hall, Inc.

– Anions (–)

• gain e-

• larger

Ionic Radius CONT’D

Ionic Radius CONT’D

• Trends within periods:– Decrease as you

move left to right

• Trends within groups:– Increase as you move – down a group

1

2

3

4 5

6

7

3. Ionization Energy

• Defined as the energy required to remove an electron from a gaseous atom

• A high ionization value indicates that the atom has a strong hold on its electrons therefore, tend to not form positive ions

• Trends within periods:– Increase left to right

• Trends within groups:– Decrease down a group

• OCTET RULE:– States that atoms tend to lose or gain electrons in

order to achieve a set of 8 valence electrons

1

2

3

4 5

6

7

4. Electronegativity

• Indicates the relative ability of atoms to attract electrons in a chemical bond

• Calculated based on many factors and are expresses in terms of a value of 4.0 or less

• Units are called Paulings• Fluorine is the most electronegative with a

value of 3.98 and Francium is the least electronegative with a value of 0.70.

Electronegativity CONT’D

• The greater the electronegativity, the more strongly it attracts the bond’s electrons

• Trends within periods:– Increases left to right

• Trends within groups:– Decreases down the

group

1

2

3

4 5

6

7