Chapter 6The Periodic Table and Periodic Law

I. History of the Periodic Table A. Just a list

• In the late 1790s, there were only 23 known elements

• The advent of electricity made it possible to break down compounds into their component elements

http://www.youtube.com/watch?v=OTEX38bQ-2w

• The development of the spectrometer made it possible to identify new elements

• By 1870 there were 70 known elements• Scientists needed a tool for organizing the

many facts associated with the elements

B. John Newlands• Noticed that when the elements were

arranged by increasing atomic mass, their properties repeated every eighth element

• A pattern such as this is called periodic

C. Mendeleev and Moseley• Mendeleev organized

the elements into the first periodic table

• He predicted the existence and properties of undiscovered elements

• Moseley rearranged the elements by increasing mass number, or number of protons, for a more accurate table

• The repetition of chemical and physical properties of elements by increasing atomic number is called the periodic law

II. The Modern Periodic Table• Elements are arranged in order of

increasing atomic number into a series of columns, called groups or families, and rows, called periods

• Each group is number 1 through 8, followed by the letter A or B

- the groups designated with an A are often referred to as the main group, or representative elements

- the groups designated with a B are referred to as the transition elements

A. Classifying the elements1. Metals• Generally shiny when smooth and clean• Solid at room temperature• Good conductors of heat and electricity• Ductile, or easily pounded into thin

sheets• Malleable, or easily drawn into wires

a. Alkali metals• Group 1A elements (except for Hydrogen)• Extremely reactive

b. Alkali earth metals• Group 2A elements• Less reactive than the alkali metals

c. Transition metalsd. Inner transition metals i. Lanthanide series

• Used extensively as phosphors, substances that emit light when struck by electrons

ii. Actinide series

2. Nonmetals• Generally gases or brittle, dull-looking

solids• Poor conductors of heat and electricity• The only liquid nonmetal is bromine (Br)

a. Halogens• Group 7A• Highly reactive

b. Noble gases• Group 8A• Extremely unreactive

3. Metalloids, or semimetals• Elements with physical and chemical

properties of both metals and nonmetals

III. Organizing the Elements by Electron Configuration

A. Valence electrons• Electrons in the highest principal energy

level• Atoms in the same group have similar

chemical properties because they have the same number of valence electrons

B. Valence electrons and period• The energy level of an element’s valence

electrons indicates the period on the periodic table in which it is found

ex. Lithium’s valence electron is in the second energy level and lithium is found in period 2

C. Valence electrons and group number• A representative element’s group number

indicates the number of valence electrons it has (with the exception of helium)

IV. The s-, p-, d-, and f-block Elements• The periodic table is divided into four

blocks

A. s-block elements• Consists of groups 1A and

2A as well as helium• Valence electrons occupy

only s orbitals• Group 1A elements have

partially filled s orbitals containing one valence electron (s1)

• Group 2A elements have completely filled s orbitals containing two valence electrons (s2)

• Because s orbitals hold a maximum of two electrons, the s-block portion of the periodic table spans two groups

B. p-block elements• After the s-block is

full, valence electrons next occupy the p sublevel and its three p orbitals

• The p-block spans six groups on the periodic table because three p orbitals can hold a maximum of six electrons

• The group 8A elements have s and p sublevels that are completely filled

- this results in an unusually stable atomic structure leaving the elements virtually unreactive

C. d-block elements• Contains the transition metals• The five d orbitals can hold a total of ten

elements; thus the d-block spans ten groups on the periodic table

• Have a filled outermost s orbital of energy level n

• Have filled or partially filled d orbitals of energy level n-1

ex. Titanium = [Ar] 4s23d2s-orbital = n = 4d-orbital = n-1 = 3

D. f-block elements• Contains the inner transition metals• Because there are seven f orbitals holding

up to a maximum of 14 electrons, the f-block spans 14 columns

E. Putting it all together• As you proceed down through the periods, the

principle energy level increases as well as the number of energy sublevels

Period 1: s-block elementsPeriods 3 and 4: s- and p-block elementsPeriods 4 and 5: s-, p-, and d-block elementsPeriods 6 and 7: s-, p-, d-, and f-block elements

V. Atomic Radius• The outer limit

of an electron cloud is defined as the spherical surface within which there is a 90% probability of finding an electron

• Atomic size is defined by how closely an atom lies to a neighboring atom, which varies from element to element

A. Trends within periods• Atomic size decreases left-to-right across a

period• Each successive element has one additional

electron in the same principal energy level• Each element also gains one more proton

that pulls the outermost electrons closer to the nucleus

B. Trends within groups• Atomic size increase as you move down a

group• A principal energy level is added to elements

in each period going down the table• Each orbital also increases in size, separating

valence electrons further from the nucleus which offsets the pull of the increased nuclear charge

Atomic Radius

VI. Ionic Radius• An ion is an atom that has a positive or

negative chargeA. When atoms lose electrons and form

positively charged ions, they always become smaller

• The electron lost is always a valence electron 1. may leave an empty outer orbital 2. the repulsion between electrons is less,

allowing them to be pulled closer to the nucleus

B. When atoms gain electrons and form negatively charged ions, they always become larger• The electrostatic repulsion between the

atom’s outer electrons forces them to move apart

C. Trends within periods 1. Positive ions on the left side of the table

become smaller left-to-right 2. Negative ions on the right side of the table

become smaller left to right

D. Trends within groups• As you move down a group, an ion’s outer

electrons are in higher principal energy levels, resulting in a gradual increase in ionic size

VII. Ionization Energy• Ionization energy is the energy required to

remove an electron from an atom• Energy is needed to overcome the

attraction between the positive charge in the nucleus and the negative charge in the electrons- a high ionization energy indicated the

atom has a strong hold on its electrons- a low ionization energy indicates an

atom loses its outer electrons easily

• Group 1A elements have low ionization energies

• Group 8A elements have high ionization energies

• The energy required to remove the first electron from an atom is called the first ionization energy

• The energy required to remove the second electron is called the second ionization energy

• For each element there is an ionization for which the required energy jumps dramatically

- related to the atom’s number of valence electrons

- atom’s hold onto their core electrons much more strongly than they hold onto their valence electrons

A. Trends within periods• Generally increases left-to-right

B. Trends within groups• Generally decreases down a group• Valence electrons farther from the nucleus

require less energy to remove them

C. Octet rule• States that atoms tend to gain, lose, or

share electrons in order to acquire a full set of eight valence electrons

- hydrogen and helium are complete with two valence electrons

1. Elements on the right side of the table tend to gain electrons (become negative ions) to acquire the noble gas configuration

2. Elements on the left side of the table tend to lose electrons (become positive ions)

VIII. Electronegativity• Indicates the relative ability of its atoms

to attract electrons in a chemical bond• Expressed in Paulings; numerical values

of 4.0 or less• Fluorine is the most electronegative

element (3.98)• Cesium (0.79) and Francium (0.7) are

the least electronegative elements

A. Trends within periods and groups• Generally decreases as you move down a

group• Increases from left-to-right