Periodic Table: Mendeleev and Modern Tables

Mendeleev's Periodic Table

• The physical and chemical properties of elements are periodic functions of their atomic weights

• Mendeleev's table has six horizontal rows (periods) and eight vertical columns (groups)

• The periodic table is based on atomic weights

• Expansion on Mendeleev's design resulted in modern periodic table, having 7 periods and 18 groups

Modern Periodic Table

• Based on atomic number, attributed to Henry Moseley in 1913
• Elements are classified based on electronic configurations

Characteristics of Periods

• Modern Periodic Table contains 7 periods
• Each period starts with an alkali metal and ends with an inert gas, with the exception of the first period
• The number of electrons in the outermost shell increases from 1 to 8 from left to right (excluding the first period)
• Elements in the third period (groups 1-2 and 13-17) are typical elements with characteristic family properties
• Valency with respect to hydrogen and chlorine increases from 1 to 4, then decreases to zero

Characteristics of Groups

• Elements in groups 1, 2, and 13-17 are normal, typical, or representative elements
• Elements in groups 3-12 are transition elements
• Elements have identical electronic configurations
• Elements in group 18 are noble gases

Advantages of the Modern Periodic Table

• Based on atomic numbers, a fundamental property of elements
• Correlates element positions with electronic configurations more clearly
• Allows prediction of new element properties before discovery
• Elements are arranged in an easy-to-remember and reproducible manner

Disadvantages of the Modern Periodic Table

• Hydrogen's position is not fixed
• Lanthanoids and Actinoids are not in the main body
• Helium is misplaced; it should be in the s-block but is in the p-block

Description

• Group 1 elements are alkali metals
• Group 2 elements are alkaline earth metals
• Group 11 elements are coinage metals
• Group 17 elements are halogens
• Group 18 (or zero group) elements are noble gases

Periodic Trends

• Major periodic trends are: atomic size/radius, metallic/non-metallic character, ionization potential energy, electron affinity, electronegativity, melting/boiling point

Atomic Size

• Atomic radius refers to the distance between the nucleus' center and the outermost shell
• Atomic size increases down a group due to added shells
• Atomic size decreases across a period because the nuclear charge increases while the number of shells remains constant, pulling electrons inward

Metallic Character

• Metals are elements which lose electrons to form positive ions (cations)
• Metallic character increases down a group because larger atomic size allows for easier electron loss
• Metallic character decreases across a period due to increased nuclear charge, making electron loss difficult

Non-Metallic Character

• Non-metals are elements which gain electrons to form negative ions (anions)
• Non-metallic character increases across a period due to increasing nuclear charge and decreasing atomic size
• Non-metallic character decreases down a group due to increasing atomic size and decreasing nuclear charge

Ionization Potential

• Ionization potential is energy needed to remove an electron from the outermost shell
• Ionization potential increases across a period because atomic size decreases as nuclear charge increases
• Ionization potential decreases down a group because atomic size increases while nuclear charge decreases

Electron Affinity

• Electron affinity is the energy change when an electron is added to an atom, forming a negative ion
• Electron affinity increases across a period due to decreasing atomic size and increasing nuclear charge
• Electron affinity decreases down a group due to increasing atomic size and decreasing nuclear charge

Electronegativity

• Electronegativity is an atom's tendency to attract a shared pair of electrons
• Electronegativity increases across a period due to increased nuclear charge
• Alkali metals have the lowest electronegativity, while halogens have the highest
• Electronegativity decreases down a group as atomic size increases and nuclear charge decreases

Melting Point Trends

• Melting point is the energy to change an element from solid to liquid
• A higher melting point indicates stronger bonds between atoms
• Melting points increase across a period and decrease down a group

Boiling Point

• Boiling point depends on the heat to transition from liquid to gas
• Boiling points increase across a period and decrease down a group

Alkali Metals

• Group 1 elements are alkali metals

• Examples are Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Caesium (Cs), and Francium (Fr)

• Francium is radioactive

• All alkali metals have one valence electron

• Alkali metals lose their valence electron and form positive ions

• Alkali metals are highly reactive

• Hydrogen, though a non-metal, is in group one (alkali metal) because it can lose electrons and form positive ions

Noble Gases

• Group 18 elements are noble/inert gases

• Examples include Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), and Radon (Rd)

• Radon is radioactive

• Noble gases are rare because of their low atmospheric abundance

• All noble gases have 8 valence electrons (octet configuration), except Helium, which has 2 (duplet configuration)

• Duplet configuration means having two electrons in the outermost shell

• Examples are Hydrogen, Helium, Lithium

• Octet configuration means having eight electrons in the outermost shell

• Examples are Helium, Neon, Argon, Krypton, Xenon, Radon

• Generally non-reactive and exist in free state