Periodic Properties and Elements

Periodic Properties and Variations

• Definitions of periodic properties and their trends in groups and periods should be studied.
• Study atomic size, metallic and non-metallic character, ionization potential, electron affinity, and electronegativity.
• Periodicity is based on atomic number.
• Understand the relation between atomic number and mass for light elements.
• Focus on the modern periodic table up to period 3 (Argon).
• Periodicity and related properties are described in terms of shells.
• Special reference to alkali metals and halogen groups.
• IUPAC recommends numbering groups from 1 to 18, but both old and new notations are acceptable.

Introduction to Elements

• Elements are pure substances with one type of atom and are the units of matter.
• Elements need classification for organized study.
• Dobereiner grouped elements in triads.
• Newland stated that every eighth element resembles the first in properties when arranged by atomic mass.
• Dmitri Mendeleev created the first periodic table based on the periodic function of atomic masses.
  • Enabled him to position 63 known elements into vertical groups and horizontal periods.
  • Though this allowed predictions of undiscovered elements, it could not explain rare earth metals or isotopes.
• Henry Moseley resolved these issues with the modern periodic table.
  • Physical and chemical properties of elements are periodic functions of their atomic number.
• Niels Bohr developed the extended form, known as the long form of the periodic table.
• A periodic table arranges elements in vertical columns (groups) and horizontal rows (periods).

Modern Periodic Table Features

• The modern periodic table consists of 18 vertical columns called groups, arranged from left to right.
• Elements are classified with atomic numbers 110 and above, that are not yet fully authenticated.
• As of 2005, the table contains 116 confirmed chemical elements; 94 occur naturally, while the rest are synthetic.
• Group 1 elements are alkali metals, which form strong alkalis with water.
• Group 2 elements are alkaline earth metals, forming weaker alkalis compared to Group 1.
• Groups 3-12 are transition elements with two incomplete outermost shells.
• Group 13 is the Boron family, with Boron as its first member.
• Group 14 is the Carbon family.
• Group 15 is the Nitrogen family.
• Group 16, the Oxygen family, is also known as chalcogens, meaning ore-forming.
• Group 17 elements are halogens, forming salts.
• Group 18 elements are noble gases or inert gases, with complete outermost orbits, making them unreactive.
• Elements in groups 1, 2, 13, 14, 15, 16, and 17 are main, representative, or normal elements.
• The outermost shells of these groups are incomplete.
• Periods are the seven horizontal rows in the modern periodic table.
• An atom's number of shells determines its period.
  • Period one elements have one shell, period two have two shells, and so on.
• Lanthanides (sixth period) and actinides (seventh period), both in Group 3, share similar properties and are placed at the bottom to prevent distortion of the table.
• Third-period elements (Na, Mg, Al, Si, P, S, and Cl) are typical elements and summarize the properties of their groups.
• A period is determined by the number of shells, while a group is determined by the number of electrons in the outermost shell.

Periodicity

• Properties that reappear or gradually vary at regular intervals are called periodic properties.
• The phenomenon is element periodicity.
• Periodicity is caused by the similar recurrence of electronic configurations of elements.
• In a group, elements have the same number of electrons in their outermost shell, so their electronic configuration is similar.
• Elements in the same group share similar chemical properties because chemical properties are determined by the number of electrons in the outer shell.

Shells and Valency

• Electrons revolve around the nucleus in orbits or shells.
• The number of shells increases down a group, equaling the period number of the element.
• Across a period, the number of shells remains constant.
• Valency is the combining capacity of an atom, or the number of electrons it can donate, accept, or share.
• Valency remains the same down a group because the number of valence electrons is constant.
• The number of valence electrons increases across a period up to Group 14, where valency is 4. After group 14, it then decreases.
• Valency is always positive, since it indicates combining capacity.

Periodic Properties

• Periodic properties are directly or indirectly related to electronic configurations and show gradation across a period or down a group.
• Important periodic properties include atomic size, metallic character, non-metallic, ionization potential, electron affinity, and electronegativity.

Atomic Size

• Atomic size / radius, is the distance between the nucleus and outermost shell.
• Alternatively, it is half the inter-nuclear distance between combined atoms in a molecule.
• Atomic size depends on the number of shells and nuclear charge.
• Adding shells increases the size of the atom.
• Increasing nuclear charge reduces the size of the atom.
• The size of an atom increases down a group due to added shells.
• In Group 1, hydrogen is the smallest.
• Across a period, the size of an atom decreases due to the increasing nuclear charge.
• Lithium has the largest atomic size in the second period, while fluorine has the smallest.
• Inert gases atoms' sizes are an exception, because they're larger because of electronic repulsion from complete outer shells.
• Cations are always smaller than their parent atoms.
• Anions are larger than their parent atoms.

Metallic and Non-Metallic Character

• Metallic elements tend to lose valence electrons to form positive ions.
• Metallic character depends on atomic size and nuclear charge.
• Greater atomic size makes it easier to remove electrons from the valence shell.
• Greater the nuclear charge makes it harder to remove electrons from the outermost orbit, decreasing metallic character.
• Elements losing electrons to complete their octet are reducing agents; metals are good reducing agents.
• Metallic nature increases down a group.
• Metallic nature decreases across a period.
• Non-metallic elements tend to gain electrons to attain an octet in their outermost orbit and are considered non-metals.
• Nonmetals typically have 5, 6, or 7 electrons in their outermost orbits and form an anion.
• Non-metallic character depends on atomic size and nuclear charge.
• Smaller the atomic size, the element is more non-metallic.
• Greater is the nuclear charge, the element is more non-metallic.
• Non-metallic nature decreases down the Group.
• Non-metallic character increases across a period.
• Non-metals are oxidizing agents.

Chemical Reactivity, Gradation and more

• Reactivity depends on the tendency to lose or gain electrons.
• Greater the tendency to lose electrons, the greater the chemical reactivity of metals.
• Greater the teendency to gain electrons, the greater the reactivity of non-metals.
• Chemical reactivity of elements first decreases and then increases across a period and the reactivity increases on going down the group.
• Melting and boiling points of metals decrease down a group, but the melting and boiling points of non-metals increases on going down the group.
• Across a period, melting and boiling points increase up to group 14 and then decrease.
• Density of elements rises gradually to maximum, with slight decreases across a period, and density rises gradually down a group.
• Group 17 (halogens) have higher electron affinity.
• Group 1 have lowest ionization energy.
• Group 18 (inert/noble gases) do not have a tendency to attract electrons.
• As atomic size increases, electronegativity decreases.
• Metals are electropositive, non-metals are electronegative.
• Elements of the second period differ in properties from their respective groups because of their high electronegativity and small atom size.
• Electronegativity increases across the period and deceases down the group. Electronegativity measures an atoms tendency to attract a shared pair of electrons towards itself known as electronegativity.
• Atomic number denotes the protons in the nuecleus, its unique to each element. Mass number is from the number of protons and neutrons in the nucleus.

Ionisation

• Electrons in an atom are attracted to the positively charged nucleus and this takes energy to overcome.
• Energy is taken to remove a neutral isolated gaseous atom, to convert it to positively charged gaseous ion which is referred to as ionisation potential (I.P.) or ionisation energy (I.E.) or first ionisation energy (IE₁).
• Unit for measurement of I.E is in electron volts per atom (eV/atom) and its S.I. unit is kilojoule per mole (kJ mol-1).
• Factor that affects the ionisation energy, is the size of the atom, the greater the size, lowers the force of attraction to the nucleus and the nuclear charge.