Atomic Radius, Ionisation Energy, Electronegativity

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

Periodicity

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. Elements often derive their names from places or people associated with their discovery:

Element 113, Nihonium, was named by Japanese researchers, as "Nihon" is the Japanese name for Japan.
Element 115, Moscovium, and element 117, Tennessine, are named after Moscow, Russia, and Tennessee, USA, respectively, in recognition of the scientists and institutions that contributed to their discovery.

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. Group 18, for instance, is made up of the noble gases, a set of colorless, odorless, and largely unreactive elements. Conversely, Group 1 houses the alkali metals, a collection of highly reactive elements that engage in vigorous reactions with water. Elements in the s and p blocks are grouped according to the number of valence electrons in their outer shell. Therefore, elements in Group 1 possess a single valence electron, conferring them with a valency of +1, and so forth.

The periodic table is split into four blocks: s, p, d, and f. These categories are assigned based on the sub-shell that accommodates the final electron in an element's atomic structure. This classification harkens back to our exploration of the Schrödinger model and the SPDF electron configuration notation, serving as a handy tool to categorise elements based on their highest energy subshell configuration.

The s and p blocks house what we typically refer to as the 'main group elements.'
The d block is reserved for elements commonly known as 'transition metals.'
The f block encompasses the lanthanides and actinides.

This classification system provides an efficient way to understand the shared properties and behaviors of the elements in the periodic table

Block	Elements	Highest energy Subshell Configurations
s-block	Groups 1 and 2 and helium	s¹ or s²
p-block	Groups 13-18 (except helium)	s²p¹ or s²p⁵
d-block	Groups 3-12	d¹s² to d¹⁰s²
f-block	Lanthanoids and actinoids	4f – subshell progressively being filled in the lanthanoids 5f – subshell progressively being filled in the actinoids

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 Year 11 Preliminary HSC syllabus, refers to these reoccurring trends observed within the periodic table.

Major periodic trends include atomic radius, state of matter at room temperature, melting point, metallic character, electronegativity, and ionisation energy. These individual trends collectively influence an element's reactivity. Emerging from the unique arrangement of the periodic table, these trends provide chemists with an indispensable tool to quickly predict an element's properties. These patterns arise due to the similar atomic structures of the elements within their respective groups or periods and the inherent periodic nature of these elements.

The diagram below offers a simplified visualisation of these trends as they appear in the periodic table

Atomic Radii Trends

The atomic size, or atomic radius, of elements gradually decreases as we move 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, its effect outweighs that of the added electrons, leading to a stronger electrostatic attraction between the positively charged protons and the negatively charged electrons. This increased attraction draws the electron shells closer to the nucleus.

In contrast, as we move down a group, the atomic radius increases. According to the Bohr atomic model, a new electron shell is added each time we descend a group, positioning the valence electrons at higher energy levels, farther from the nucleus. Although the electron count rises, the electrostatic attraction between the nucleus and the electrons doesn't significantly increase. This is attributed to the "electron shielding effect," where electrons in the preceding shells create a repelling force, detracting from the positive pull on the new valence electrons. As a result of this shielding effect, valence electrons are more loosely held.

Core Charge

The core charge of an atom represents the effective nuclear charge felt by a valence electron. By calculating an atom's core charge, we can predict the properties of different elements more accurately and explain the observed trends on the periodic table.

Because of the electron shielding effect discussed earlier, an element's core charge is calculated as the total number of protons minus the number of inner shell electrons.

State at Room Temperature

The state of elements at room temperature generally shifts from solid to less solid, or even gaseous, as we transition from left to right across a period. Elements within the S and D blocks are typically solid, while elements in the later P block groups, such as groups 16, 17, and 18, are often gaseous.

All elements in group 18, known as 'noble gases', exist as gases at room temperature. Notable exceptions to these trends include mercury in the D block and bromine in group 17 of the P block, both of which are liquid at room temperature.

Melting Point

The melting point refers to the specific temperature at which a substance transitions from its solid to liquid state. This change occurs when supplied heat energy disrupts the bonds linking an element's atoms. Substances with stronger bonds necessitate more energy to break apart these atomic connections. Melting points across the periodic table vary significantly and don't conform to a particular trend.

Metallic Character

The metallic character of an element reflects how easily an atom can relinquish a valence electron. The metallic character heightens as we move from right to left across a period since the pull between the valence electrons and the nucleus diminishes, facilitating electron loss.

As we descend a group, the metallic character also increases due to the expansion of the atomic size. When the atomic size swells, the outer electron shells are located further from the nucleus and, therefore, experience less attraction, making it easier for these shells to lose electrons

Electronegativity

lectronegativity refers to an atom's ability to attract a pair of electrons for bond formation. This chemical property is influenced by the atom's atomic number and the distance that separates its valence electrons from the nucleus. Many atoms abide by the octet rule, seeking to hold eight electrons in their valence shell.

As we move left to right across the periodic table, electronegativity tends to increase. Elements on the left have less than half-filled valence shells, making it energetically costly to gain electrons as compared to losing them. Consequently, these elements often lose electrons when forming bonds. In contrast, elements on the right side can more efficiently gain electrons to complete their outer valence shell, thereby becoming more electronegative.
When we go top to bottom down a group, electronegativity tends to decrease. With the addition of more electrons and corresponding energy levels, the atomic radius expands. As a result, the attraction between the nucleus and the valence electrons diminishes due to the increased distance.
Elements in the D and F blocks, including transition metals, lanthanides, and actinides, do not adhere to these trends. Their metallic properties impact their ability to attract electrons, while noble gases, with their complete valence shells, do not attract additional electrons.

Ionisation Energy

Ionisation energy, denoted as Ei, signifies the energy required to detach the most loosely bound electron (the valence electron) from a neutral atom in its gaseous phase. Essentially, ionisation is the reverse of electronegativity, reflecting an atom's tendency to release an electron and form an ion rather than attract a pair of electrons.

Ions are charged atoms that form when an electron is either added to or removed from an atom's valence shell. Positively charged ions are called cations, while negatively charged ones are known as anions.

A higher ionisation energy suggests that an atom is less likely to become a cation, as more energy is needed to remove the valence electron. Elements on the right-hand side of the periodic table typically have higher ionisation energies due to their nearly filled valence shells.

In general, ionisation energy within a period tends to increase from left to right. As anticipated, noble gases have very high ionisation energies, thanks to their fully occupied valence shells.