Classification of Elements and Periodicity in Properties - Notes

Classification of Elements and Periodicity in Properties

Objectives

Appreciate how grouping elements based on properties led to the Periodic Table.
Understand the Periodic Law.
Understand the significance of atomic number and electronic configuration for periodic classification.
Name elements with Z > 100 using IUPAC nomenclature.
Classify elements into s, p, d, f blocks and learn their characteristics.
Recognize periodic trends in physical and chemical properties.
Compare element reactivity and correlate with their natural occurrence.
Explain the relationship between ionization enthalpy and metallic character.
Use scientific vocabulary for atomic/ionic radii, ionization enthalpy, electron gain enthalpy, electronegativity, and valence.

Introduction

The Periodic Table is a crucial concept in chemistry for students, researchers, and professionals.
It demonstrates that elements are not randomly assorted but show trends and form families.
Understanding the Periodic Table is essential for comprehending the fundamental building blocks of chemistry.

2.1 Why Classify Elements?

Elements are the basic units of matter.
In 1800, 31 elements were known; by 1865, the number doubled to 63; currently, 114 elements are known, with ongoing efforts to synthesize more.
Studying the chemistry of each element and their compounds individually is challenging due to the large number of elements.
Scientists classify elements to organize knowledge, rationalize chemical facts, and predict new ones.

2.2 Genesis of Periodic Classification

Classification and the Periodic Law result from systematizing knowledge gained through observations and experiments.
Johann Dobereiner (early 1800s) identified trends among element properties, noting similarities within groups of three elements (Triads) by 1829.
In Triads, the middle element's atomic weight was approximately halfway between the other two, and its properties were intermediate.
- Dobereiner's Triads:
  - Lithium (Li): Atomic weight 7
  - Sodium (Na): Atomic weight 23
  - Potassium (K): Atomic weight 39
  - Calcium (Ca): Atomic weight 40
  - Strontium (Sr): Atomic weight 88
  - Barium (Ba): Atomic weight 137
  - Chlorine (Cl): Atomic weight 35.5
  - Bromine (Br): Atomic weight 80
  - Iodine (I): Atomic weight 127
The Law of Triads was dismissed as coincidence because it only worked for a few elements.
A.E.B. de Chancourtois (1862) arranged elements by increasing atomic weights in a cylindrical table, but this received little attention.
John Alexander Newlands (1865) proposed the Law of Octaves, arranging elements by atomic weights and noting that every eighth element had similar properties -- similar to musical octaves.
- Newlands' Octaves:
  - Li (7), Be (9), B (11), C (12), N (14), O (16), F (19)
  - Na (23), Mg (24), Al (27), Si (29), P (31), S (32), Cl (35.5)
  - K (39), Ca (40)
The Law of Octaves was only true for elements up to calcium.
Newlands was later awarded the Davy Medal in 1887 by the Royal Society, London, for his work.
Dmitri Mendeleev (1834-1907) and Lothar Meyer (1830-1895) independently developed the Periodic Law.
In 1869, both chemists proposed that elements arranged by increasing atomic weights show regular intervals in physical and chemical properties.
Lothar Meyer plotted physical properties like atomic volume, melting point, and boiling point against atomic weight, observing a periodically repeated pattern with changes in the length of the repeating pattern.
By 1868, Meyer had developed a table resembling the Modern Periodic Table, but his work was published after Mendeleev's.
Mendeleev is credited with first publishing the Periodic Law:
- "The properties of the elements are a periodic function of their atomic weights."
Mendeleev arranged elements in rows and columns by increasing atomic weights, placing elements with similar properties in the same column or group.
Mendeleev's classification was more elaborate than Meyer's and recognized the significance of periodicity, using a broader range of physical and chemical properties.
Mendeleev relied on similarities in empirical formulas and compound properties.
He sometimes ignored atomic weight order to group similar elements together, assuming atomic weight measurements were incorrect.
- For example, iodine (lower atomic weight) was placed with fluorine, chlorine, and bromine due to similar properties, despite tellurium having a higher atomic weight.
Mendeleev left gaps in the table, predicting undiscovered elements such as gallium and germanium, which he called Eka-Aluminium and Eka-Silicon, respectively.
Mendeleev predicted the properties of these elements, which were later discovered.
- Mendeleev’s Predictions for Eka-aluminium (Gallium) and Eka-silicon (Germanium):
  - Eka-aluminium (predicted):
    - Atomic weight: 68
    - Density: 5.9 g/cm³
    - Melting point: Low
    - Formula of oxide: $E<em>2O</em>3$
    - Formula of chloride: $ECl_3$
  - Gallium (found):
    - Atomic weight: 70
    - Density: 5.94 g/cm³
    - Melting point: 302.93 K
    - Formula of oxide: $Ga<em>2O</em>3$
    - Formula of chloride: $GaCl_3$
  - Eka-silicon (predicted):
    - Atomic weight: 72
    - Density: 5.5 g/cm³
    - Melting point: High
    - Formula of oxide: $EO_2$
    - Formula of chloride: $ECl_4$
  - Germanium (found):
    - Atomic weight: 72.6
    - Density: 5.36 g/cm³
    - Melting point: High
    - Formula of oxide: $GeO_2$
    - Formula of chloride: $GeCl_4$
Mendeleev's quantitative predictions and their accuracy made him and his Periodic Table famous.

2.3 Modern Periodic Law and the Present Form of the Periodic Table

When Mendeleev developed his table, the internal structure of the atom was unknown.
In 1913, Henry Moseley observed regularities in the X-ray spectra of elements.
Plotting $\sqrt{v}$ (where $v$ is the frequency of X-rays emitted) against atomic number (Z) yielded a straight line, unlike plots against atomic mass.
Moseley demonstrated that atomic number is a more fundamental element property than atomic mass.
The Modern Periodic Law states:
- "The physical and chemical properties of the elements are periodic functions of their atomic numbers."
The Periodic Law revealed analogies among the 94 naturally occurring elements and stimulated interest in inorganic chemistry, leading to the creation of artificial elements.
Atomic number equals nuclear charge (number of protons) or the number of electrons in a neutral atom.
The Periodic Law is a consequence of periodic variation in electronic configurations, which determine physical and chemical properties.
Numerous Periodic Table forms exist, some emphasizing chemical reactions and valence, others focusing on electronic configuration.
The "long form" of the Periodic Table is most convenient and widely used.
Horizontal rows are called periods, and vertical columns are called groups.
Elements with similar outer electronic configurations are arranged in vertical columns (groups or families).
IUPAC recommends numbering groups from 1 to 18, replacing older notations.
There are seven periods; the period number corresponds to the highest principal quantum number (n).
The first period has 2 elements; subsequent periods have 8, 8, 18, 18, and 32 elements, respectively.
The seventh period is incomplete but theoretically has a maximum of 32 elements.
14 elements of the sixth and seventh periods (lanthanoids and actinoids) are placed in separate panels at the bottom.

2.4 Nomenclature of Elements with Atomic Numbers > 100

Traditionally, the discoverer names new elements, subject to IUPAC ratification.
Controversies have arisen because new elements with high atomic numbers are highly unstable and available in minute quantities.
Synthesis and characterization require sophisticated equipment, leading to competition among laboratories.
Scientists may prematurely claim discovery before collecting reliable data.
For example, both American and Soviet scientists claimed discovery of element 104, naming it Rutherfordium and Kurchatovium, respectively.
IUPAC recommends a systematic nomenclature derived from the atomic number using numerical roots until discovery is proven and the name is officially recognized.
Roots for digits 0-9 are:
- 0 = nil (n)
- 1 = un (u)
- 2 = bi (b)
- 3 = tri (t)
- 4 = quad (q)
- 5 = pent (p)
- 6 = hex (h)
- 7 = sept (s)
- 8 = oct (o)
- 9 = enn (e)
The roots are combined in the order of the digits in the atomic number, and "ium" is added at the end.
IUPAC names for elements with Z > 100 are listed (see table in the transcript).
A new element initially receives a temporary name and symbol consisting of three letters.
A permanent name and symbol are later assigned by a vote of IUPAC representatives, reflecting the discovery country or honoring a notable scientist.
As of now, elements up to 112, 114, and 116 have been discovered; 113, 115, 117, and 118 are not yet known.
Problem: What is the IUPAC name and symbol for element 120?
- Solution: unbinilium (Ubn)

2.5 Electronic Configurations of Elements and the Periodic Table

An electron in an atom is characterized by four quantum numbers; the principal quantum number (n) defines the main energy level or shell.
Electrons fill different subshells or orbitals (s, p, d, f) in an atom, resulting in its electronic configuration.
An element's position in the Periodic Table reflects the quantum numbers of the last orbital filled.

(a) Electronic Configurations in Periods

The period indicates the value of n for the outermost or valence shell.
Successive periods correspond to filling the next higher principal energy level (n = 1, n = 2, etc.).
The number of elements in each period is twice the number of atomic orbitals available in the energy level being filled.
The first period (n = 1) starts with filling the 1s level, with hydrogen ( $1s^1$ ) and helium ( $1s^2$ ) completing the K shell.
The second period (n = 2) starts with lithium, and the third electron enters the 2s orbital; beryllium has the configuration $1s^2 2s^2$ .
From boron, the 2p orbitals fill until neon, which completes the L shell ( $2s^2 2p^6$ ). Thus, there are 8 elements in the second period.
The third period (n = 3) begins with sodium, where the added electron enters the 3s orbital. Successive filling of 3s and 3p orbitals results in 8 elements, from sodium to argon.
The fourth period (n = 4) starts with potassium, and the added electrons fill the 4s orbital. Before the 4p orbital fills, the 3d orbitals become energetically favorable, leading to the 3d transition series.
Scandium (Z = 21) has the electronic configuration $3d^1 4s^2$ . The 3d orbitals are filled at zinc (Z = 30), with the configuration $3d^{10} 4s^2$ .
The fourth period ends at krypton with the filling of the 4p orbitals, totaling 18 elements.
The fifth period (n = 5) begins with rubidium and is similar to the fourth period, containing the 4d transition series starting at yttrium (Z = 39).
This period ends at xenon with the filling of the 5p orbitals.
The sixth period (n = 6) contains 32 elements, with electrons successively entering 6s, 4f, 5d, and 6p orbitals.
Filling of the 4f orbitals begins with cerium (Z = 58) and ends at lutetium (Z = 71), forming the 4f inner transition series or lanthanide series.
The seventh period (n = 7) is similar to the sixth period, with successive filling of the 7s, 5f, 6d, and 7p orbitals, including most man-made radioactive elements.
This period will end at element 118, a noble gas. Filling the 5f orbitals after actinium (Z = 89) gives the 5f inner transition series or actinide series.
The 4f and 5f inner transition series are placed separately to maintain structure and group elements with similar properties.
Problem: How would you justify the presence of 18 elements in the 5th period of the Periodic Table?
- Solution: For n = 5, l = 0, 1, 2, 3. The order of increasing energy is 5s < 4d < 5p. There are 9 available orbitals, accommodating 18 electrons, thus 18 elements.

(b) Groupwise Electronic Configurations

Elements in the same vertical column (group) have similar valence shell electronic configurations, the same number of outer electrons, and similar properties.
Group 1 elements (alkali metals) all have $ns^1$ valence shell electronic configurations.
- Li: [ $\text{He}$ ] $2s^1$
- Na: [ $\text{Ne}$ ] $3s^1$
- K: [ $\text{Ar}$ ] $4s^1$
- Rb: [ $\text{Kr}$ ] $5s^1$
- Cs: [ $\text{Xe}$ ] $6s^1$
- Fr: [ $\text{Rn}$ ] $7s^1$
Element properties have periodic dependence on atomic number, not relative atomic mass.

2.6 Electronic Configurations and Types of Elements: s, p, d, f Blocks

The Aufbau principle and electronic configurations give theoretical foundation for periodic classification.
Elements in a vertical column constitute a group or family, exhibiting similar chemical behavior due to the same number and distribution of electrons in their outermost orbitals.
Elements are classified into s-block, p-block, d-block, and f-block based on the type of atomic orbital being filled with the last electron.
Exceptions include helium, which belongs to the s-block but is placed in the p-block due to its filled valence shell ( $1s^2$ ), exhibiting noble gas properties. Hydrogen, with one electron, can be placed in Group 1 or Group 17, but is placed separately at the top of the table because it is a special case.

2.6.1 The s-Block Elements

Group 1 (alkali metals) and Group 2 (alkaline earth metals) have $ns^1$ and $ns^2$ outermost electronic configurations, respectively.
They are reactive metals with low ionization enthalpies, readily losing outermost electrons to form 1+ or 2+ ions.
Metallic character and reactivity increase down the group.
Due to high reactivity, they are never found pure in nature.
Compounds of s-block elements, except those of lithium and beryllium, are predominantly ionic.

2.6.2 The p-Block Elements

p-Block elements comprise Groups 13 to 18; together with s-block elements, they are called Representative Elements or Main Group Elements.
Outermost electronic configurations vary from $ns^2np^1$ to $ns^2np^6$ in each period.
Each period ends with a noble gas element having a closed valence shell $ns^2np^6$ configuration.
Noble gases have low chemical reactivity due to the difficulty of altering their stable configuration.
Halogens (Group 17) and chalcogens (Group 16) are chemically important non-metals with highly negative electron gain enthalpies, readily adding one or two electrons to attain a stable noble gas configuration.
Non-metallic character increases from left to right across a period, while metallic character increases down the group.

2.6.3 The d-Block Elements (Transition Elements)

d-Block elements are in Groups 3 to 12 in the center of the Periodic Table, characterized by filling inner d orbitals with electrons.
They have the general outer electronic configuration $(n-1)d^{1-10} ns^{1-2}$ .
They are all metals, mostly forming colored ions, exhibiting variable valence (oxidation states), paramagnetism, and are often used as catalysts.
Zn, Cd, and Hg, with configuration $(n-1)d^{10} ns^2$ , do not show most transition element properties.
Transition metals bridge chemically active s-block metals and less active elements of Groups 13 and 14.

2.6.4 The f-Block Elements (Inner-Transition Elements)

The two rows at the bottom of the Periodic Table are Lanthanoids (Ce(Z = 58) - Lu(Z = 71)) and Actinoids (Th(Z = 90) - Lr(Z = 103)).
They have the outer electronic configuration $(n-2)f^{1-14} (n-1)d^{0-1}ns^2$ , with the last electron entering the f-orbital.
These are called Inner-Transition Elements (f-Block Elements) and are all metals.
Elements within each series have similar properties.
Early actinoid chemistry is more complicated than lanthanoid chemistry due to the large number of possible oxidation states.
Actinoid elements are radioactive.
Many actinoid elements are made in nanogram quantities or less by nuclear reactions, and their chemistry is not fully studied.
Elements after uranium are called Transuranium Elements.
Problem: The elements Z = 117 and 120 have not yet been discovered. In which family / group would you place these elements and also give the electronic configuration in each case.
- Solution: Element Z = 117 belongs to the halogen family (Group 17) with electronic configuration [ $\text{Rn}$ ] $5f^{14}6d^{10}7s^27p^5$ . Element Z = 120 will be placed in Group 2 (alkaline earth metals) with electronic configuration [ $\text{Uuo}$ ] $8s^2$ .