The Periodic Table and Periodic Law

Chapter 6

1. History of the Periodic Table’s Development

• In the 1700s, Lavoisier compiled a list of all the known elements of the time.

• List contained 23 elements

• The 1800s brought large amounts of information, including an explosion in the number of known elements.

• Approximately 70 known elements

• Scientists needed a way to organize knowledge about elements.

• A significant step came when chemists agreed upon a method for accurately determining the atomic masses of the elements.

Section 6.1 - Development of the Modern Periodic Table

History of the Periodic Table’s Development (cont.)

• John Newlands• In 1864, English chemist John

Newlands (1837-1898) proposed an arrangement where elements were ordered by increasing atomic mass.

• Newlands noticed when the elements were arranged by increasing atomic mass, their properties repeated every eighth element.

• The pattern was called periodic because it repeats in a specific manner.

• He called it the law of octaves

• Unfortunately the law didn’t work for all the known elements

History of the Periodic Table’s Development (cont.)

• Meyer, Mendeleev, & Moseley• In 1869, German chemist Lothar Meyer and

Russian chemist Dmitri Mendeleev both demonstrated a connection between atomic mass and elemental properties.

• Mendeleev arranged the table by increasing atomic mass, which resulted in a periodic pattern in properties, the first periodic table.

• He predicted the existence and properties of undiscovered elements.

• He left blank spaces where he thought undiscovered elements should go

• He predicted the properties of scandium, gallium, and germanium

History of the Periodic Table’s Development (cont.)

• Mendeleev’s table was not completely correct, some elements were not in the correct order

• In 1913, English chemist Henry Moseley discovered that atoms of each element contain unique number of protons in their nuclei.

• rearranged the table by increasing atomic number, and resulted in a clear periodic pattern.

• Periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number is called periodic law.

2. The Modern Periodic Table

• The modern periodic table contains boxes which contain the element's name, symbol, atomic number, and atomic mass.

• Columns of elements are called groups.

• Rows of elements are called periods.

• Elements in groups 1,2, and 13-18 possess a wide variety of chemical and physical properties and are called the representative elements.

• Elements in groups 3-12 are known as the transition metals.

The Modern Periodic Table (cont.)

• Elements are classified as metals,

non-metals, and metalloids.

• Metals• Metals are elements that are

generally shiny when smooth and clean, solid at room temperature, and good conductors of heat and electricity.

• Alkali metals are all the elements in group 1 except hydrogen, and are very reactive.

• Alkaline earth metals are in group 2, and are also highly reactive.

The Modern Periodic Table (cont.)

• The transition elements are divided into transition metals and inner transition metals.

• The two sets of inner transition metals are called the lanthanide series and actinide series and are located at the bottom of the periodic table.

The Modern Periodic Table (cont.)

• Non-metals

• Non-metals are elements that are generally gases or brittle, dull-looking solids, and poor conductors of heat and electricity.

• Group 17 is composed of highly reactive elements called halogens.

• Group 18 gases are extremely unreactive and commonly called noble gases.

The Modern Periodic Table (cont.)

• Metalloids

• Metalloids have physical and chemical properties of both metals and non-metals.

End of Section 6.1

3. Organizing the Elements by Electron Configuration

• Recall electrons in the highest principal energy level are called valence electrons.

• All group 1 elements have one valence electron.

• The energy level of an element’s valence electrons indicates the period on the periodic table in which it is found.

• The number of valence electrons for elements in groups 13-18 is ten less than their group number.

Section 6.2 - Classification of the Elements

4. The s-, p-, d-, and f-Block Elements

• The shape of the periodic table becomes clear if it is divided into blocks representing the atom’s energy sublevel being filled with valence electrons.

• s-block elements consist of groups 1 and 2, and the element helium.

• Group 1 elements have a partially filled s orbital with one electron.

• Group 2 elements have a completely filled s orbital with two electrons.

• After the s-orbital is filled, valence electrons occupy the p-orbital.

• Groups 13-18 contain elements with completely or partially filled p orbitals.

The s-, p-, d-, and f-Block Elements (cont.)

• The d-block contains the transition metals and is the largest block.

• There are exceptions, but d-block elements usually have filled outermost s orbital, and filled or partially filled d orbital.

• The five d orbitals can hold 10 electrons, so the d-block spans ten groups on the periodic table.

• The f-block contains the inner transition metals.

• f-block elements have filled or partially filled outermost s orbitals and filled or partially filled 4f and 5f orbitals.

• The 7 f orbitals hold 14 electrons, and the inner transition metals span 14 groups.

End of Section 6.2

5. Atomic Radius • Many properties of the elements tend to change in a predictable way, known as a trend, as you move across a period or down a group.

• Atomic size is a periodic trend defined by how closely an atom lies to a neighboring atom.

• The electron cloud does not have a clearly defined edge since it is based on probability.

• For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element.

• For elements that occur as molecules (nonmetals), the atomic radius is half the distance between nuclei of identical atoms.

Section 6.3 – Periodic Trends

Atomic Radius (cont.)

• Trends Within Periods• There is a general decrease in

atomic radii from left to right across a period.

• Caused by increasing positive charge in the nucleus and the fact that the principal energy level remains the same.

• Valence electrons are not shielded from the increasing nuclear charge because no additional electrons come between the nucleus and the valence electrons.

• Thus, the increased nuclear charge pulls the outermost electrons closer to the nucleus.

Atomic Radius (cont.)

• Trends Within Groups• Atomic radii generally increases as

you move down a group.• The nuclear charge increases as

electrons are added to successively higher principal energy levels.

• The outermost orbital size increases down a group with energy level, making the atom larger.

• Therefore, outer electrons are farther from the nucleus offsetting the increase nuclear charge.

• Also, inner electrons shield outer electrons from the pull of the nucleus.

6. Ionic Radius

• An ion is an atom or bonded group of atoms with a positive or negative charge.

• When atoms lose electrons and form positively charged ions, they always become smaller for two reasons:

1.The loss of a valence electron can leave an empty outer orbital resulting in a small radius.

2.Electrostatic repulsion decreases allowing the electrons to be pulled closer to the radius.

• When atoms gain electrons, they can become larger, because the addition of an electron increases electrostatic repulsion.

Ionic Radius (cont.)

• Trends Within Periods• The ionic radii of positive ions

generally decrease from left to right.

• The ionic radii of negative ions generally decrease from left to right, beginning with group 15 or 16.

• Trends Within Groups• In general, there is a gradual

increase in ionic size as you move down a group.

• Both positive and negative ions increase in size moving down a group.

7. Ionization Energy • Ionization energy is defined as the energy required to remove an electron from a gaseous atom.

• The energy required to remove the first electron is called the first ionization energy.

• Removing the second electron requires more energy, and is called the second ionization energy.

• Each successive ionization requires more energy, but it is not a steady increase.

Ionization Energy (cont.)

• A high ionization energy value indicates the atom has a strong hold on its electrons, and therefore less likely to form positive ions.

• Likewise, a low ionization energy value indicates an atom loses its outer electrons easily, are likely to form positive ions.

• Trends Within Periods• First ionization energy generally increases as

you move left to right across a period.• Increased nuclear charge produces an

increased hold on valence electrons.

• Trends Within Groups• First ionization energy generally decreases as

you move down the group.• With electrons farther away from the nucleus,

less energy required to remove them.

Ionization Energy (cont.)

• Octet Rule• The octet rule states that atoms tend to

gain, lose, or share electrons in order to acquire a full set of eight valence electrons.

8. Electronegativity

• The electronegativity of an element indicates the relative ability of its atoms to attract electrons in a chemical bond.

• Electronegativity generally decreases as you move down a group, and increases as you move left to right across a period.