Introduction to The Periodic Table & Atomic Radius

This is part of preliminary HSC Chemistry course under the topic of Periodicity

HSC Chemistry Syllabus

  • Demonstrate, explain, and predict the relationships in the observable trends in the physical and chemical properties of elements in periods and groups in the periodic table, including but not limited to:
– State of matter at room temperature
– Electron configurations and atomic radii
– First ionisation energy and electronegativity
– Reactivity with water

      Periodic Table: The Basics

      The Modern Periodic Table 

      The modern periodic table, illustrated above, expands upon Mendeleev's original categorisation of 63 elements by arranging the elements according to increasing atomic number rather than atomic mass. Despite these modifications, the fundamental structure of rows and columns, reminiscent of Mendeleev's design, is preserved as it reflects the inherent order of the elements.

       

       

      Periods

      The rows in the modern periodic table are referred to as periods, echoing the nomenclature used in Mendeleev's original table. As we move from left to right across a period, each successive element has one more proton than its predecessor.

      The length of periods in the modern periodic table can vary: the first period comprises only hydrogen and helium, while periods 6 and 7 are extended, necessitating that many of their elements be placed beneath the main body of the table.

      In Bohr's model of the atom, electrons occupy shells (orbits). The period in which an element is found also represents the number of electron shells its atom contains. For example, hydrogen atom, found in the first period, has 1 electron shell. Sodium atom, found in the third period, has 3 electron shells.

      Groups

      The columns of the periodic table are known as groups. The modern table hosts 18 groups, with elements within the same group generally sharing similar properties.

      For example, group 18 is made up of the noble gases, a set of colourless, odourless, and largely unreactive elements. Conversely, Group 1 houses the alkali metals, a collection of highly reactive elements that engage in vigorous reactions with water.

       

      Valence electrons, valency and oxidation state

       

      While periods reveal the number of electron shells an atom has, groups reveal the number of valence electrons in the atom (with exception of transition metals). For example, sodium, in group 1, has 1 valence electron; calcium in group 2, has 2 valence electrons; and chlorine in group 17, has 7 valence electrons.

       

      The valence or valency refers to the typical number of bonds an atom of a given element forms when it's part of a molecule of compound. 

      For example, 

      • Carbon (group 14) has a valency of 4 which means an atom of carbon typically forms four bonds with other atoms in a molecule
      • Chlorine (group 17) has a valency of 1 which means an atom of chlorine typically forms one bond with another atom in a molecule

       

      The oxidation state refers to the number of electrons needed to be either gained or lost to complete the outermost (valence) shell. A positive valence suggests the element needs to lose electrons (negatively charged), and a negative valence suggests the element needs to gain electrons. In contrast to the valency number, the oxidation state can be positive (for an electropositive atom) or negative (for an electronegative atom).

      For example:

      • Aluminium (group 13) has an oxidation state of +3 which means it needs to lose 3 electrons in its valence shell to achieve a complete outermost shell (octet rule)
      • Chlorine (group 17) has an oxidation state of –1 which means it needs to gain 1 electron in its valence shell to achieve a complete outermost shell
      • Noble gases (group 18) have an oxidation state of 0 because their atoms already contain a complete valence shell

      Many transition metals have more than one typical oxidation state, and thus they are not shown in the periodic table above.

      Metals, Non-metals, and Metalloids

       

      Metals are located on the left (exception of hydrogen) and in the centre of the periodic table. The structure of pure metals consists of the same atoms arranged in layers. The electrons of these atoms are delocalised instead of fixed in an atom's orbit.

      • Group 1 metals are known as alkali metals
      • Group 2 metals are known as alkali Earth metals
      • Group 3 to 12 are transition metals

       

      Non-metals are found on the right side of the periodic table. The structure of non-metals varies between covalent molecular and covalent lattice network. This is discussed separately here.

      • Group 17 elements are known as halogens
      • Group 18 elements are known as noble gases

       

      Metalloids (semi-metals) are situated along the zig-zag line separating metals and nonmetals, metalloids exhibit properties intermediate between those of metals and nonmetals and include elements like silicon and arsenic.

      What Are Periodic Table Trends?

      Periodic trends are specific patterns that emerge in the periodic table, illustrating various characteristics of elements, including their size and electronic properties. The term "periodicity," as mentioned in Module 1: Properties and Structure of Matter in the HSC syllabus, refers to these reoccurring trends observed within the periodic table including:

      • Atomic and ionic radii

      Atomic Radii Trends

      The atomic radius refers to the physical size of an atom of a particular element.

       

       

      Atomic radius gradually decreases from left to right across a period. This decrease results from the simultaneous addition of electrons and protons to the atom. As the proton number increases, the effective nuclear charge of the atom increases, leading to a stronger electrostatic attraction between the positively charged protons and the negatively charged electrons. This increased attraction brings the electron shells closer to the nucleus, resulting in a reduced atomic radius. While additional electrons in an atom results in increased electrostatic repulsion between newly added and existing electrons, this is outweighed by the additional attractive force caused by the protons in the nucleus.

       

      In contrast, atomic radius increases down a group. As we descend down a group, a few things change:

      • the number of electron shells increases; this increases the atomic radius as a new shell is further away from the nucleus.
      • the number of protons and electrons increases; this decreases the atomic radius due to greater electrostatic attraction between the nucleus and electron shells.

      While the two changes result in opposite effects on the atomic radius, the addition of electron shells usually outweighs the effect of more protons and electrons. This is due to the electron shielding effect, where electrons in the preceding shells create a repelling force on the electrons in the newly added shell, reducing the effect of attraction exerted by the protons in the nucleus. In simpler words, electron shielding decreases the effective nuclear charge. 

      Atomic vs Ionic Radii

      The difference in radius between an atom and its ions (both cations and anions) stems from changes in the electron configuration and the effective nuclear charge experienced by the electrons. Sodium and chlorine serve as classic examples to illustrate these differences due to their common ionic forms, Na⁺ and Cl⁻, respectively.

       

      Cations generally have smaller radii compared to their atoms; anions generally have larger radii compared to their atoms.

       

      Sodium (Na)

      • Atomic Radius of Sodium: The atomic radius of a sodium atom is determined by the distance from the nucleus to the outermost electron in its valence shell. Sodium has one valence electron.
      • Formation of Sodium Ion (Na⁺): When sodium forms a cation (Na⁺), it loses its one valence electron, resulting in a reduction of the electron-electron repulsion in the ion and leaving the sodium ion with a complete octet in the next lower shell. This loss of an electron also decreases the number of electron shells, which also contributes to a smaller radius.
      • As a result of these changes, the radius of a Na⁺ ion is smaller than the atomic radius of a neutral sodium atom. The removal of the outermost electron and the increased effective nuclear charge cause the electron shells to contract.

      Chlorine (Cl)

      • Atomic Radius of Chlorine: The atomic radius of a chlorine atom is defined by the distance from the nucleus to the outermost electrons in its valence shell. Chlorine has seven valence electrons.
      • Formation of Chloride Ion (Cl⁻): Chlorine forms an anion (Cl⁻) by gaining an electron to complete its valence shell. This addition of an electron increases electron-electron repulsion within the ion while keeping the number of electron shells unchanged.
      • The radius of a Cl⁻ ion is larger than the atomic radius of a neutral chlorine atom. The increase in electron-electron repulsion with the addition of an extra electron outweighs the increase in effective nuclear charge, causing the electron shells to expand.

       

      RETURN TO MODULE 1: PROPERTIES AND STRUCTURE OF MATTER