Detailed Notes on Classification of Elements and Periodicity in Properties
The Periodic Table
The Periodic Table is a fundamental concept in chemistry, providing organization and revealing trends among elements.
It's essential for students, researchers, and anyone seeking to understand the building blocks of the chemical world.
Glenn T. Seaborg emphasized its importance in understanding the fundamental building blocks of chemistry.
Unit 3: Classification of Elements and Periodicity in Properties
Learning Objectives:
Understand the historical development and significance of the Periodic Table.
Comprehend the Periodic Law and its basis in atomic number and electronic configuration.
Learn to name elements with atomic numbers greater than 100 using IUPAC nomenclature.
Classify elements into s, p, d, and f blocks based on their electronic configurations.
Recognize periodic trends in physical and chemical properties of elements.
Compare element reactivity and relate it to their occurrence in nature.
Explain the relationship between ionization enthalpy and metallic character.
Use scientific vocabulary to communicate ideas related to atomic properties (e.g., atomic/ionic radii, ionization enthalpy, electronegativity).
Topics Covered:
Historical development of the Periodic Table
Modern Periodic Law
Periodic classification based on electronic configuration
Periodic trends in physical and chemical properties
3.1 Why Classify Elements?
Elements are the basic units of matter.
The number of known elements has increased significantly over time: 31 in 1800, 63 in 1865, and 114 at present.
Classifying elements systematically helps to:
Organize knowledge of elements and their compounds.
Rationalize known chemical facts.
Predict new elements and their properties.
3.2 Genesis of Periodic Classification
The Periodic Law and Table resulted from systematizing knowledge from various scientists' observations and experiments.
Dobereiner's Triads (Early 1800s)
Johann Dobereiner noted similarities in physical and chemical properties among groups of three elements (Triads).
The middle element's atomic weight was approximately halfway between the other two.
Properties of the middle element were intermediate to the other two.
Example Triads:
Lithium (Li), Sodium (Na), Potassium (K)
Li = 7 Na = 23 K = 39
Calcium (Ca), Strontium (Sr), Barium (Ba)
Ca = 40 Sr = 88 Ba = 137
Chlorine (Cl), Bromine (Br), Iodine (I)
Cl = 35.5 Br = 80 I = 127
The Law of Triads was limited and dismissed as coincidence.
De Chancourtois' Cylindrical Table (1862)
A.E.B. de Chancourtois arranged elements by increasing atomic weights in a cylindrical table.
This displayed periodic recurrence of properties.
The approach did not gain much attention.
Newlands' Law of Octaves (1865)
John Alexander Newlands arranged elements in increasing order of atomic weights.
Every eighth element had properties similar to the first, like octaves in music.
Example Octaves:
Li, Be, B, C, N, O, F
Li = 7 Be = 9 B = 11 C = 12 N = 14 O = 16 F = 19
Na, Mg, Al, Si, P, S, Cl
Na = 23 Mg = 24 Al = 27 Si = 29 P = 31 S = 32 Cl = 35.5
K, Ca
K = 39 Ca = 40
The law was true only for elements up to calcium.
Newlands was later awarded the Davy Medal by the Royal Society, London, in 1887.
Mendeleev and Meyer's Contributions (1869)
Dmitri Mendeleev and Lothar Meyer independently proposed that arranging elements by increasing atomic weights revealed similarities in physical and chemical properties at regular intervals.
Lothar Meyer plotted physical properties (atomic volume, melting/boiling point) against atomic weight, observing a periodically repeated pattern.
Meyer noted changes in the length of the repeating pattern, and by 1868, had a table resembling the Modern Periodic Table.
Mendeleev is generally credited with developing the Modern Periodic Table.
Mendeleev's Periodic Law
Mendeleev published the Periodic Law for the first time:
"The properties of the elements are a periodic function of their atomic weights."
He arranged elements in horizontal rows and vertical columns by increasing atomic weights.
Elements with similar properties occupied the same vertical column (group).
Mendeleev's system was more elaborate than Meyer's and recognized the significance of periodicity.
He used a broader range of physical and chemical properties (especially empirical formulas and compound properties) for classification.
Mendeleev sometimes ignored the strict order of atomic weights to group elements with similar properties (believing atomic measurements might be incorrect).
Example: Iodine (lower atomic weight) placed with fluorine, chlorine, and bromine due to similar properties.
He left gaps in the table for undiscovered elements and predicted their properties.
Examples: Eka-Aluminium (Gallium) and Eka-Silicon (Germanium).
Predicted properties such as atomic weight, density, melting point, oxide formula, and chloride formula.
Mendeleev's Predictions vs. Actual Properties
Eka-aluminium (predicted) vs. Gallium (found):
Atomic weight: 68 vs. 70
Density: 5.9 g/cm3 vs. 5.94 g/cm3
Melting point: Low vs. 302.93 K
Formula of oxide: E2O3 vs. Ga2O3
Formula of chloride: ECl3 vs. GaCl3
Eka-silicon (predicted) vs. Germanium (found):
Atomic weight: 72 vs. 72.6
Density: 5.5 g/cm3 vs. 5.36 g/cm3
Melting point: High vs. 1231 K
Formula of oxide: EO2 vs. GeO2
Formula of chloride: ECl4 vs. GeCl4
The success of Mendeleev's predictions made him and his Periodic Table famous.
3.3 Modern Periodic Law and the Present Form of the Periodic Table
Mendeleev developed his table without knowledge of the internal structure of atoms.
Moseley's Contribution (1913)
Henry Moseley observed regularities in the X-ray spectra of elements.
Plotting \sqrt{\nu} (frequency of X-rays) against atomic number (Z) gave a straight line, unlike plots against atomic mass.
He showed that atomic number is a more fundamental property than atomic mass.
Modern Periodic Law
The Modern Periodic Law states:
"The physical and chemical properties of the elements are periodic functions of their atomic numbers."
The law revealed analogies among the 94 naturally occurring elements and stimulated interest in inorganic chemistry.
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.
Electronic configurations determine the physical and chemical properties of elements and their compounds.
Modern Periodic Table
Various forms of the Periodic Table exist.
The "long form" is the most convenient and widely used.
Horizontal rows are called periods, and vertical columns are called groups or families.
Elements with similar outer electronic configurations are arranged in vertical columns.
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) of the elements.
The first period contains 2 elements, and subsequent periods contain 8, 8, 18, 18, and 32 elements, respectively.
The seventh period is incomplete but theoretically has a maximum of 32 elements.
Lanthanoids and actinoids (14 elements each) are placed in separate panels at the bottom to maintain the table's structure.
3.4 Nomenclature of Elements with Atomic Numbers > 100
Traditionally, discoverers named new elements with IUPAC ratification.
Controversies arose due to the instability and minute quantities of new elements.
The IUPAC recommends a systematic nomenclature derived directly from the atomic number using numerical roots.
IUPAC Nomenclature
Numerical roots for digits 0-9:
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)
Roots are combined in order of digits, and "ium" is added at the end.
The temporary symbol consists of three letters.
A permanent name and symbol are determined by IUPAC representatives.
Official names of elements up to 118 have been announced by IUPAC.
Examples of IUPAC Names and Symbols
101: Unnilunium (Unu) - Mendelevium (Md)
102: Unnilbium (Unb) - Nobelium (No)
103: Unniltrium (Unt) - Lawrencium (Lr)
104: Unnilquadium (Unq) - Rutherfordium (Rf)
105: Unnilpentium (Unp) - Dubnium (Db)
106: Unnilhexium (Unh) - Seaborgium (Sg)
107: Unnilseptium (Uns) - Bohrium (Bh)
108: Unniloctium (Uno) - Hassium (Hs)
109: Unnilennium (Une) - Meitnerium (Mt)
110: Ununnillium (Uun) - Darmstadtium (Ds)
111: Unununnium (Uuu) - Roentgenium (Rg)
112: Ununbium (Uub) - Copernicium (Cn)
113: Ununtrium (Uut) - Nihonium (Nh)
114: Ununquadium (Uuq) - Flerovium (Fl)
115: Ununpentium (Uup) - Moscovium (Mc)
116: Ununhexium (Uuh) - Livermorium (Lv)
117: Ununseptium (Uus) - Tennessine (Ts)
118: Ununoctium (Uuo) - Oganesson (Og)
Problem 3.1
What is the IUPAC name and symbol for the element with atomic number 120?
Solution: The roots for 1, 2, and 0 are un, bi, and nil, respectively.
Therefore, the symbol is Ubn, and the name is unbinilium.
3.5 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 (shell).
The distribution of electrons into orbitals is called electronic configuration.
An element's location 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.
Period 1 (n=1)
Starts with filling the 1s level and has two elements:
Hydrogen (1s^1)
Helium (1s^2)
The first shell (K) is completed.
Period 2 (n=2)
Starts with lithium, and the third electron enters the 2s orbital.
Lithium: 1s^22s^1
Beryllium has four electrons (1s^22s^2).
From Boron, the 2p orbitals are filled until Neon (2s^22p^6).
There are 8 elements in the second period.
Period 3 (n=3)
Begins at Sodium, with the added electron entering a 3s orbital.
Successive filling of 3s and 3p orbitals gives rise to the third period of 8 elements from sodium to argon.
Period 4 (n=4)
Starts at Potassium, and the added electrons fill up the 4s orbital.
Before the 4p orbital is filled, the filling up of 3d orbitals becomes energetically favorable.
This gives rise to the 3d transition series of elements.
Scandium (Z=21) : 3d^14s^2
The 3d orbitals are filled at Zinc (Z=30) with electronic configuration 3d^{10}4s^2
The fourth period ends at Krypton with the filling up of the 4p orbitals.
Altogether there are 18 elements in this fourth period.
Period 5 (n=5)
Beginning with Rubidium, it is similar to the fourth period and contains the 4d transition series starting at Yttrium (Z=39).
This period ends at Xenon with the filling up of the 5p orbitals.
Period 6 (n=6)
Contains 32 elements and successive electrons enter 6s, 4f, 5d, and 6p orbitals.
The filling of the 4f orbitals begins with Cerium (Z=58) and ends at Lutetium (Z=71) to give the 4f-inner transition series called the lanthanoid series.
Period 7 (n=7)
Similar to the sixth period but includes most man-made radioactive elements.
Electrons successively fill the 7s, 5f, 6d, and 7p orbitals.
This period will end at the element with atomic number 118, which would belong to the noble gas family.
Filling of the 5f orbitals after actinium (Z = 89) gives the 5f-inner transition series known as the actinoid series.
The 4f- and 5f-inner transition series of elements are placed separately in the Periodic Table to maintain its structure.
Problem 3.2
How would you justify the presence of 18 elements in the 5th period of the Periodic Table?
Solution: When n = 5, l = 0, 1, 2, 3. The order in which the energy of the available orbitals 4d, 5s, and 5p increases is 5s < 4d < 5p.
The total number of orbitals available is 9. The maximum number of electrons that can be accommodated is 18; therefore, 18 elements are in the 5th period.
(b) Groupwise Electronic Configurations
Elements in the same vertical column or group have similar valence shell electronic configurations, the same number of electrons in the outer orbitals, and similar properties.
Example: Group 1 (Alkali Metals)
All have ns^1 valence shell electronic configuration.
Lithium (Li): [He]2s^1
Sodium (Na): [Ne]3s^1
Potassium (K): [Ar]4s^1
Rubidium (Rb): [Kr]5s^1
Cesium (Cs): [Xe]6s^1
Francium (Fr): [Rn]7s^1
3.6 Electronic Configurations and Types of Elements: s-, p-, d-, f- Blocks
The aufbau principle and electronic configuration of atoms provide a theoretical foundation for the periodic classification.
Elements in a vertical column of the Periodic Table constitute a group or family and exhibit similar chemical behavior.
This similarity arises because these elements have the same number and distribution of electrons in their outermost orbitals.
The elements can be classified into four blocks (s-block, p-block, d-block, and f-block) depending on the type of atomic orbitals that are being filled with electrons.
Exceptions to Block Categorization
Helium belongs to the s-block but is positioned in the p-block because it has a completely filled valence shell (1s^2) and exhibits properties characteristic of noble gases.
Hydrogen has only one s-electron and can be placed in group 1 (alkali metals) or group 17 (halogens).
Because it is a special case, hydrogen is placed separately at the top of the Periodic Table.
3.6.1 The s-Block Elements
Comprise Groups 1 and 2, having ns^1 and ns^2 outermost electronic configurations.
Reactive metals with low ionization enthalpies.
Readily lose outermost electron(s) to form 1+ or 2+ ions.
Metallic character and reactivity increase down the group.
Never found pure in nature due to high reactivity.
Compounds are predominantly ionic, except for those of lithium and beryllium.
3.6.2 The p-Block Elements
Comprise Groups 13 to 18.
Outermost electronic configuration varies from ns^2np^1 to ns^2np^6.
Representive Elements or Main Group Elements together with s-block Elements.
At the end of each period is a noble gas element with a closed valence shell (ns^2np^6) configuration.
Noble gases exhibit very low chemical reactivity.
Halogens (Group 17) and chalcogens (Group 16) are chemically important non-metals with highly negative electron gain enthalpies.
Non-metallic character increases across a period; metallic character increases down the group.
3.6.3 The d-Block Elements (Transition Elements)
Comprise Groups 3 to 12 in the center of the Periodic Table.
Characterized by filling inner d orbitals by electrons.
General outer electronic configuration is (n-1)d^{1-10}ns^{0-2}, except for Pd (4d^{10}5s^0).
All metals.
Mostly form colored ions, exhibit variable valence (oxidation states), paramagnetism, and are used as catalysts.
Zn, Cd, and Hg ((n-1)d^{10}ns^2) do not show typical transition element properties.
Form a bridge between chemically active metals of s-block elements and less active elements of Groups 13 and 14.
3.6.4 The f-Block Elements (Inner-Transition Elements)
Two rows at the bottom of the Periodic Table: Lanthanoids (Ce to Lu) and Actinoids (Th to Lr).
Outer electronic configuration (n-2)f^{1-14}(n-1)d^{0-1}ns^2.
The last electron added to each element is filled in the f-orbital.
All metals.
Properties of elements within each series are similar.
The chemistry of early actinoids is more complicated due to a larger number of possible oxidation states.
Actinoid elements are radioactive.
Many actinoid elements exist in nanogram quantities and their chemistry remains under investigation.
Elements after uranium are called Transuranium Elements.
Problem 3.3
The elements Z = 117 and 120 have not yet been discovered. In which family/group would you place these elements, and give the electronic configuration in each case?
Solution:
Z = 117 would belong to the halogen family (Group 17) with electronic configuration [Rn] 5f^{14}6d^{10}7s^27p^5
Z = 120 will be placed in Group 2 (alkaline earth metals) with electronic configuration [Uuo]8s^2
3.6.5 Metals, Non-metals, and Metalloids
Elements can be divided into Metals and Non-Metals.
Metals
Comprise more than 78% of all known elements and appear on the left side of the Periodic Table.
Usually solids at room temperature (mercury is an exception).
High melting and boiling points.
Good conductors of heat and electricity.
Malleable and ductile.
Non-metals
Located at the top right-hand side of the Periodic Table.
Usually solids or gases at room temperature with low melting and boiling points (boron and carbon are exceptions).
Poor conductors of heat and electricity.
Most non-metallic solids are brittle and are neither malleable nor ductile.
Metalloids
Elements bordering the zig-zag line (e.g., silicon, germanium, arsenic, antimony, and tellurium).
Show properties characteristic of both metals and non-metals.
Problem 3.4
Considering atomic number and position in the periodic table, arrange the following elements in increasing order of metallic character: Si, Be, Mg, Na, P.
Solution:
Metallic character increases down a group and decreases along a period from left to right.
Hence the order of increasing metallic character is: P < Si < Be < Mg < Na.
3.7 Periodic Trends in Properties of Elements
Observable patterns in physical and chemical properties as we descend in a group or move across a period.
Chemical reactivity tends to be high in Group 1 metals, lower in elements towards the middle of the table, and increases to a maximum in the Group 17 non-metals.
Reactivity increases down a group of representative metals (alkali metals) but decreases down a group of non-metals (halogens).
3.7.1 Trends in Physical Properties
Numerous physical properties show periodic variations, including melting and boiling points, heats of fusion and vaporization, and energy of atomization.
Periodic trends are discussed with respect to atomic and ionic radii, ionization enthalpy, electron gain enthalpy, and electronegativity.
(a) Atomic Radius
Finding the size of an atom is more complicated than measuring the radius of a ball.
The electron cloud surrounding the atom does not have a sharp boundary.
An estimate of the atomic size can be made by knowing the distance between atoms in the combined state.
Covalent radius: Half the distance between two atoms bound together by a single bond in a covalent molecule.
Metallic radius: Half the internuclear distance separating the metal cores in the metallic crystal.
Atomic radii can be measured by X-ray or other spectroscopic methods.
Atomic size generally decreases across a period and increases down a group.
Trends in Atomic Radius
Across a Period:
Atomic size generally decreases due to an increase in effective nuclear charge within the same valence shell.
Down a Group:
Atomic radius increases due to the increase in the principal quantum number and shielding of outer electrons from the nucleus by inner electrons.
Noble gases radii are large because their (non-bonded radii) values are very large.
Radii of noble gases should be compared not with the covalent radii but with the van der Waals radii of other elements.
(b) Ionic Radius
Removal of an electron from an atom results in a cation, whereas the gain of an electron leads to an anion.
Ionic radii can be estimated by measuring the distances between cations and anions in ionic crystals.
Ionic radii of elements exhibit the same trend as atomic radii.
Trends in Ionic Radius
Cation is smaller than its parent atom because it has fewer electrons with the same nuclear charge.
Anion is larger than its parent atom due to increased repulsion among electrons and a decrease in effective nuclear charge.
Isoelectronic species have the same number of electrons.
Radii vary due to different nuclear charges.
Greater positive charge = smaller radius.
Greater negative charge = larger radius.
(c) Ionization Enthalpy
A quantitative measure of the tendency of an element to lose electrons.
Represents the energy required to remove an electron from an isolated gaseous atom (X) in its ground state.
X(g) \rightarrow X^+(g) + e^-
Measured in kJ/mol.
Second ionization enthalpy is higher than the first.
X^+(g) \rightarrow X^{2+}(g) + e^-
Ionization enthalpies are always positive.
First ionization enthalpies generally increase across a period and decrease down a group.
Factors Affecting Ionization Enthalpy
The attraction of electrons towards the nucleus.
The repulsion of electrons from each other.
Effective nuclear charge is the net positive charge experienced by valence electrons.
Shielding or screening describes the repulsion of valence electrons by inner electrons.
Trends in Ionization Enthalpy
Across a Period:
Ionization enthalpy increases due to increasing nuclear charge outweighing the shielding effect.
Down a Group:
Ionization enthalpy decreases due to increased shielding and the outermost electron being farther from the nucleus.
Anomalies in Ionization Enthalpy
Boron (Z = 5) has lower first ionization enthalpy than beryllium (Z = 4).
Easier to remove 2p-electron from boron compared to 2s-electron from beryllium due to higher penetration of s-electrons.
Oxygen (Z = 8) has a smaller first ionization enthalpy compared to nitrogen (Z = 7).
In oxygen, two 2p-electrons occupy the same 2p-orbital, resulting in increased electron-electron repulsion, making it easier to remove an electron.
(d) Electron Gain Enthalpy
Enthalpy change when an electron is added to a neutral gaseous atom to convert it into a negative ion.
X(g) + e^- \rightarrow X^-(g)
Provides a measure of the ease with which an atom adds an electron.
Can be either endothermic or exothermic.
Halogens have very high negative electron gain enthalpies.
Trends in Electron Gain Enthalpy
Electron gain enthalpy becomes more negative with increasing atomic number across a period.
Electron gain enthalpy becomes less negative down a group.
Electron gain enthalpy of O or F is less negative than that of the succeeding element.
The added electron goes to the smaller n = 2 quantum level and suffers significant repulsion from other electrons present in this level.
(e) Electronegativity
A qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself.
Numerical scales include Pauling scale, Mulliken-Jaffe scale, and Allred-Rochow scale.
Pauling assigned a value of 4.0 to fluorine.
Electronegativity is not constant for a given element and varies depending on the element to which it is bound.
Trends in Electronegativity
Generally increases across a period from left to right.
Decreases down a group.
Attraction between outer electrons and the nucleus increases as atomic radius decreases in a period.
Inversely related to metallic properties of elements.
3.7.2 Periodic Trends in Chemical Properties
Trends such as diagonal relationships, inert pair effect, and effects of lanthanoid contraction.
(a) Periodicity of Valence or Oxidation States
Valence is the most characteristic property of elements.
Valence of representative elements is usually equal to the number of electrons in the outermost orbitals or eight minus the number of outermost electrons.
Oxidation state can be defined as the charge acquired by an atom based on electronegative consideration from other atoms in the molecule.
(b) Anomalous Properties of Second-Period Elements
The first element of groups 1, 2, and 13-17 differs in many respects from other members of their respective group.
Due to small size, large charge/radius ratio, and high electronegativity.
Only four valence orbitals (2s and 2p) available for bonding, limiting maximum covalency to 4.
Greater ability to form p\pi - p\pi multiple bonds to itself and other second-period elements.
3.7.3 Periodic Trends and Chemical Reactivity
Chemical reactivity is related to electronic configuration.
Atomic and ionic radii generally decrease across a period from left to right.
Ionization enthalpies increase and electron gain enthalpies become more negative across a period.
Chemical reactivity is highest at the two extremes and lowest in the center.
Metallic character decreases and non-metallic character increases moving from left to right across the period.
In a group, metallic character increases and non-metallic character decreases down the group.
Summary
The Periodic Law and the Periodic Table have evolved over time.
Mendeleev's Periodic Table was based on atomic mass.
The Modern Periodic Table is based on atomic number and arranges elements in periods (rows) and groups (columns).
Electronic configuration explains similar chemical properties of elements in the same group.
Elements are classified into s-block, p-block, d-block, and f-block based on their electronic configurations.
Metals, non-metals, and metalloids exhibit different properties.
Periodic trends are observed in atomic sizes, ionization enthalpies, electron gain enthalpies, electronegativity, and valence.