Classification of Elements and Periodicity in Properties — Key Vocabulary
3.1 WHY DO WE NEED TO CLASSIFY ELEMENTS ?
Elements are the basic units of all matter.
By 1800, only 31 elements were known; by 1865, identified elements had more than doubled to 63; today, 114 elements are known.
Many are man-made; efforts to synthesize new elements continue.
Problem: with so many elements, studying each one individually is impractical.
Solution: classify elements to organize knowledge, rationalize known facts, and predict new elements for future study.
Outcomes: a systematic framework to relate properties, foresee trends, and guide research.
3.2 GENESIS OF PERIODIC CLASSIFICATION
Periodic classification emerged from systematic observations by multiple scientists; it groups elements with similar properties.
Dobereiner (early 1800s): proposed Triads (groups of three elements) and noted the middle element’s atomic weight was about half-way between the other two; its properties lay between those of the other two. This suggested recurring patterns, later termed periodicity.
Example: Triads with Li–Na–K; Cl–Br–I; Ca–Sr–Ba (Table 3.1 in the text).
John A. Newlands (1865): Law of Octaves.
Arranged elements by increasing atomic weight and observed that every eighth element had properties similar to the first; analogy to musical octaves.
Effective only up to calcium; not universally accepted.
Recognition: award of the Davy Medal (1887) by the Royal Society, London, for his work.
A. E. B. de Chancourtois (1862): proposed a cylindrical (tortuous) arrangement of elements by increasing atomic weights to reveal periodicity (an early form of the 3D organization).
Lothar Meyer (German) and Dmitri Mendeleev (Russian) independently studied periodicity; Meyer plotted physical properties (atomic volume, melting/boiling points) against atomic weight and observed periodic patterns; his table resembled the modern form but slower to publish.
Mendeleev’s contributions (1869 onwards):
Published the Periodic Law: The properties of the elements are a periodic function of their atomic weights.
Classified elements into horizontal rows (periods) and vertical columns (groups) to place elements with similar properties in the same group.
Introduced predictions for undiscovered elements, leaving gaps in his table for elements yet to be found, which showcased the power of his periodic law.
Important flexibility: he occasionally ignored strict atomic-weight order to place elements with similar properties together; he left gaps for unknown elements (e.g., Eka-Aluminium and Eka-Silicon) and predicted their properties.
Predictions (Table 3.3) matched later discoveries: Gallium (Ga) and Germanium (Ge) properties aligned with Mendeleev’s predictions.
3.3 MODERN PERIODIC LAW AND THE PRESENT FORM OF THE PERIODIC TABLE
Shift from atomic mass to atomic number (Z) as the fundamental organizing parameter.
Henry Moseley (1913) linked X-ray spectral frequencies to atomic number; plotted frequency vs. atomic number (not vs atomic mass) and obtained a straight line.
Conclusion: the atomic number (proton count/electron count in a neutral atom) is the true basis of periodicity, not atomic weight.
Modern Periodic Law (as revised by Moseley): The physical and chemical properties of the elements are periodic functions of their atomic numbers.
Consequences of the modern law:
Periodic table now reflects electronic structure and the arrangement of electrons in shells and subshells.
There are seven periods and eighteen groups.
Elements in the same group have similar outer electron configurations and hence similar chemical properties.
The table has many forms; the long form is the most convenient, emphasizing periods (horizontal) and groups (vertical).
Nomenclature and structure:
IUPAC groups numbered 1–18 (replacing older IA–VIIA, VIII, IB–VIIB, 0).
Periods correspond to the principal quantum number n of the outermost electrons.
The seventh period is incomplete; the maximum theoretical length is 32, similar to the sixth period.
The long form places the lanthanoids and actinoids in separate bottom panels (below the main table) to maintain the regular pattern of groups.
3.4 NOMENCLATURE OF ELEMENTS WITH ATOMIC NUMBERS > 100
Naming of new elements historically followed the discoverers’ preference, ratified by IUPAC.
For very unstable, recently synthesized elements, systematic nomenclature is used until discovery is confirmed:
The temporary name is built from numerical roots corresponding to the digits of the atomic number, with the suffix -ium.
Examples: Z = 101–118 use roots (un, nil, etc.) with suffix -ium; e.g., 101 = Unnilunium (symbol Unu), 102 = Unnilbium (Uub), etc., leading to temporary symbols like Uuu, Ubn, etc., as shown in Table 3.4.
IUPAC-named elements (permanent names and symbols) are assigned after confirmation and voting by IUPAC representatives; some elements beyond 100 use dynamic naming conventions tied to the country, scientists, or notable figures.
Glenn T. Seaborg’s work and the placement of actinides below lanthanoids:
Seaborgium (Sg, element 106) named in honor of Seaborg.
By the time of the text, elements up to Z = 118 have been discovered and officially named.
Problem 3.1 (example): What is the IUPAC name and symbol for Z = 120?
Answer: Roots for 1, 2, and 0 are un, bi, nil respectively; symbol and name: Ubn and unbinilium.
3.5 ELECTRONIC CONFIGURATIONS OF ELEMENTS AND THE PERIODIC TABLE
Core idea: an atom’s electrons occupy orbitals (s, p, d, f) according to the Aufbau principle; the distribution is the electronic configuration.
The last (outermost) orbital filled determines the element’s position in the long form periodic table.
(a) Electronic configurations in Periods
The period number equals the value of n for the outermost shell; each successive period corresponds to filling the next principal quantum number (n = 1, 2, …).
Period lengths reflect the number of orbitals available at that energy level and how many electrons can be accommodated:
Period 1 (n = 1): 2 elements → 1s^1 and 1s^2 (H and He).
Period 2 (n = 2): 8 elements → start Li, end Ne with configurations ending in 2p^6.
Period 3 (n = 3): 8 elements → Na to Ar; 3s and 3p fill after 3s^1–3s^2 then 3p^6.
Period 4 (n = 4): 18 elements → K to Kr; 3d orbitals begin to fill as 4p fills after 3d1–3d10; transition series begin with Sc (Z = 21) 3d^1 4s^2; Zn (Z = 30) completes 3d^10 4s^2; 4p fills to Kr.
Period 5 (n = 5): 18 elements → starts with Rb; 4d series begins with Y (Z = 39) through Xe; 5p fills to end of period with Xe.
Period 6 (n = 6): 32 elements → fills 6s, 4f, 5d, 6p; 4f begins at Ce (Z = 58) and ends at Lu (Z = 71) forming the 4f lanthanoids; 5f begins with Th (Z = 90) and ends before elements above U; 6p ends with noble gas configuration at end of the period; includes lanthanide contraction considerations.
Period 7 (n = 7): similar to the sixth; includes 7s, 5f, 6d, 7p; ends at element 118 (Og).
(b) Groupwise electronic configurations
Elements in the same vertical group have similar valence shell configurations, same number of outer-shell electrons, and similar chemistry.
(c) The four blocks: s-block, p-block, d-block, f-block
s-block: Groups 1–2; outer ns^1 and ns^2; features alkali and alkaline earth metals.
p-block: Groups 13–18; outer ns^2np^1 to ns^2np^6; includes halogens and noble gases; representative elements.
d-block: Groups 3–12; filling of inner d orbitals; transition elements; metallic; variable oxidation states; often colored ions and catalysts; Zn, Cd, Hg are exceptions with full d^10 configuration.
f-block: Lanthanoids (Ce–Lu) and Actinoids (Th–Lr); outer configuration (n−2)f^1–14 (n−1)d^0–1 ns^2; inner-transition elements; all metals; many actinoids are radioactive; chemistry less explored due to small quantities.
(d) Metals, Non-metals, and Metalloids
Metals: left side, typically solids at room temperature (except Hg); high melting/boiling points; good conductors; malleable and ductile.
Non-metals: top-right; poor conductors; often solids, liquids, or gases with low melting/boiling points; brittle when solid.
Metalloids: border-line elements (Si, Ge, As, Sb, Te) with properties between metals and non-metals; shown as a zig-zag boundary in Fig. 3.3.
Problem 3.3 (examples): Z = 117 in Group 17; Z = 120 in Group 2; electronic configurations: [Rn] 5f^14 6d^10 7s^2 7p^5 for Z = 117; [Uuo] 8s^2 for Z = 120.
3.6 THE BLOCKS AND ELEMENT TYPES IN DETAIL
3.6.1 The s-Block Elements
Groups 1 and 2: alkali metals and alkaline earth metals; outer configurations ns^1 and ns^2.
General properties: highly reactive metals; reactive increases down the group; typically form +1 (alkali) or +2 (alkaline earth) cations; compounds (except Li and Be) are largely ionic.
Examples in the table: Li, Na, K, Rb, Cs, Fr with configurations starting from [He]2s^1 to [Rn]7s^1.
3.6.2 The p-Block Elements
Groups 13–18 (the Representative Elements or Main Group Elements).
Outer configurations vary from ns^2np^1 (Group 13) to ns^2np^6 (Group 18).
Noble gas end of each period; halogens (Group 17) and chalcogens (Group 16) are chemically important non-metals; general trend of increasing non-metallic character across a period; increasing metallic character down a group.
3.6.3 The d-Block Elements (Transition Elements)
Groups 3–12; typical outer configuration (n−1)d^1–10 ns^0–2; except Pd: [Kr] 4d^{10} 5s^0.
Metals; often colored ions; variable oxidation states; catalytic properties; paramagnetism; Zn, Cd, Hg (d^10) often not true transition metals by some definitions.
Act as a bridge between highly reactive s-block metals and less reactive p-block elements.
3.6.4 The f-Block Elements (Inner-Transition Elements)
Lanthanoids: Ce (Z = 58)–Lu (Z = 71).
Actinoids: Th (Z = 90)–Lr (Z = 103).
Outer configuration: (n−2)f^1–14 (n−1)d^0–1 ns^2.
The f-block elements are all metals; properties within each series are quite similar; actinoids are mainly radioactive and many are synthetically produced; many are not fully studied due to extremely small quantities.
3.6.5 Metals, Non-metals, and Metalloids
Metals constitute >78% of known elements and occupy the left and middle parts of the table.
Non-metals dominate the upper-right portion; poor conductors; more volatile/less metallic.
Metalloids (semi-metals) border metal and non-metal sides; e.g., Si, Ge, As, Sb, Te; show mixed properties and lie along the zig-zag line between metals and non-metals.
Figure references: Fig. 3.3 showing the blocks and metal/non-metal/metalloid classifications.
3.7 PERIODIC TRENDS IN PROPERTIES OF ELEMENTS
There are many periodic patterns in physical and chemical properties as you descend a group or move across a period; key properties discussed: atomic radii, ionic radii, ionization enthalpy, electron gain enthalpy, electronegativity.
(3.7.1) Trends in Physical Properties
Atomic Radius (and covalent/metallic radii):
Distance between nuclei in a bonded state or crystal provides an estimate of atomic size.
Covalent radius: half the bond length between atoms in a covalent bond (example: Cl-Cl bond length ≈ 198 pm ⇒ covalent radius ≈ 99 pm).
Metallic radius: half the internuclear distance between adjacent metal atoms in a metallic lattice (example: Cu-Cu ≈ 256 pm ⇒ metallic radius ≈ 128 pm).
General trends:
Across a period, atomic radius decreases due to increasing effective nuclear charge (Z_eff) drawing electrons closer.
Down a group, atomic radius increases due to higher principal quantum number and shielding by inner shells.
Noble gases radii are not included in the simple covalent radius trend; their nonbonded radii are large and, when compared, are better described by van der Waals radii.
Table 3.6 shows representative atomic radii values for Period II and Period III elements; e.g., Li 152 pm, Be 111 pm, B 88 pm, C 77 pm, N 74 pm, O 66 pm, F 64 pm (Period II); Na 186, Mg 160, Al 143, Si 117, P 110, S 104, Cl 99 (Period III).
Ionic Radius
Cations: smaller than parent atoms due to electron removal while nuclear charge remains; anions: larger due to added electrons and electron-electron repulsion.
Isoelectronic series (same electron count) show sizes based on nuclear charge: higher Z ⇒ smaller radius.
Example: F–, O^{2−}, Na^+, Mg^{2+} are isoelectronic; Na^+ smaller than Na, etc.
Ionization Enthalpy (First Ionization Enthalpy, ∆iH1)
Defined as energy required to remove first electron from a neutral gaseous atom: X(g) → X^+(g) + e^−; unit: kJ mol^−1.
Always positive; second ionization enthalpy ∆iH2 is the energy to remove the second electron from X^+(g): X^+(g) → X^{2+}(g) + e^−, and so on.
∆iH generally increases across a period and decreases down a group (Fig. 3.5, 3.6a, 3.6b).
Be and B anomaly: Be has higher ∆iH than B because 2s electrons are more penetrating and less shielded than 2p electrons; within the same principal quantum number, s-electrons are more tightly bound to nucleus than p-electrons.
Oxygen vs Nitrogen anomaly: O has lower ∆iH than N due to electron-electron repulsion when two electrons pair in the same 2p orbital, versus Hund’s rule in N where electrons occupy separate orbitals.
Relation to shielding: across a period, increased nuclear charge outweighs shielding; down a group, shielding increases faster than nuclear charge, reducing ∆iH.
Electron Gain Enthalpy (∆egH, sometimes called electron affinity)
Energy change when adding an electron to a neutral atom: X(g) + e^− → X^−(g).
Negative values indicate exothermic attachment (energy released); halogens have highly negative ∆egH as they gain electrons to attain noble gas configurations; noble gases have positive ∆egH due to adding an electron to a new, less stable shell.
Across a period, ∆egH generally becomes more negative; down a group, ∆egH tends to become less negative due to increasing size and poorer attraction between added electron and nucleus.
Anomalies: O and F have less negative ∆egH compared to the next element because added electrons enter the smaller 2p shell, causing repulsion; in larger shells (3p), the added electron can occupy more space with less repulsion.
Table 3.7 (sample values) shows ∆egH for H, Li, Na, K, etc., and Group 16, 17 trends.
Electronegativity
Qualitative measure of an atom’s tendency to attract shared electrons in a chemical bond; scales include Pauling, Mulliken, Allred-Rochow; Pauling scale is most widely used.
Pauling values (examples): Li 1.0, Be 1.5, B 2.0, C 2.5, N 3.0, O 3.5, F 4.0 (Period II); Na 0.9, Mg 1.2, Al 1.5, Si 1.8, P 2.1, S 2.5, Cl 3.0 (Period III); down a group, electronegativity decreases (e.g., Li > Na > K; F > Cl > Br > I).
Trend explanations: across a period, decreasing atomic radius increases nucleus–valence electron attraction, increasing electronegativity; down a group, increasing atomic radius and shielding reduce attraction, decreasing electronegativity.
(3.7.2) Periodic Trends in Chemical Properties
Valence/oxidation states show periodicity: valence often equals the number of electrons in the outermost shell or eight minus that number.
Oxidation state concept illustrated with reactions: Na2O and OF2 demonstrate the transfer of electrons dictated by electronegativity differences; oxygen often takes electrons from metals; fluorine is highly electronegative and attains −1 in many compounds like HF.
Example: SiBr4 (Si in Group 14, valence 4) and Al2S3 (Al in Group 13, valence 3; S in Group 16, valence 2).
Periodic trends of valence and oxidation states show diagonal relationships (e.g., Li–Mg and Be–Al) and explain some deviations in the early members of groups (anomalous behaviour of Li and Be).
Diagonal relationships: small size and high electronegativity of the first members (Li, Be, B, etc.) lead to covalent character and limitations on maximum covalency.
Problem 3.4 ordering metallic character: P < Si < Be < Mg < Na (metallic character increases down a group and decreases across a period).
(3.7.3) Periodic Trends and Chemical Reactivity
Reactivity is highest at the extremes of a period: leftmost (alkali metals) can lose electrons readily (low ∆iH); rightmost (halogens) readily gain electrons (high ∆egH negative).
Center elements have moderate reactivity; metals tend to form oxides that are basic (e.g., Na2O), non-metals form acidic oxides (e.g., Cl2O7); some oxides (Al2O3, As2O3) are amphoteric.
For transition elements, changes in atomic radii are smaller across a period; ionization enthalpies are intermediate between s- and p-blocks; they are less electropositive than group 1 and 2 metals.
Overall, metallic character increases down a group and non-metallic character decreases down a group for main group elements; trends for transition elements can be opposite due to d-electron involvement.
SUMMARY
The Periodic Law and Periodic Table emerged from efforts to classify elements and understand repeating patterns.
Modern Periodic Law (Moseley) ties properties to atomic number, not atomic mass, aligning with electronic structure.
The long form Periodic Table organizes elements into 7 periods and 18 groups, with s-, p-, d-, and f-blocks reflecting the type of subshell being filled.
Elements are categorized as metals, non-metals, or metalloids; their properties reflect their electronic structure and position in the table.
Periodic trends (atomic/ionic radii, ionization enthalpy, electron gain enthalpy, electronegativity) arise from a balance of nuclear charge and electron shielding; explain reactivity and bonding tendencies.
Valence and oxidation states follow periodic rules but include anomalies (diagonal relationships, inert pair effect, lanthanoid contraction) that are important in predicting chemistry.
Nomenclature for new elements follows IUPAC rules, with temporary names based on digits and eventual official names after confirmation.
Problem sets (3.1–3.40) reinforce how to apply periodic trends, electron configurations, and nomenclature to real examples.
Problems and Illustrative Solutions (selected)
Problem 3.1: What would be the IUPAC name and symbol for Z = 120?
Answer: Unbinilium, symbol Ubn.
Problem 3.3: Where would Z = 117 and Z = 120 fit in the table, and what are the suggested electron configurations?
Z = 117: Group 17 (halogens); electronic configuration [Rn] 5f^14 6d^10 7s^2 7p^5.
Z = 120: Group 2 (alkaline earth metals); electronic configuration [Uuo] 8s^2.
Problem 3.4: Why does the 6th period contain 32 elements?
Explanation: When n = 5 (as in the 5th period), the available orbitals are 5s, 4d, and 5p with total 9 orbitals and 18 electrons possible; for n = 6, the available orbitals include 6s, 4f, 5d, and 6p; accounting for all orbitals gives 32 electrons, hence 32 elements in the 6th period.
Problem 3.8: Using the Periodic Table, predict formulas of compounds: (a) Si with Br -> SiBr4; (b) Al with S -> Al2S3.
Problem 3.15: Energy of electron in ground state of hydrogen is -2.18×10^-18 J. Calculate ∆H ionization for hydrogen in kJ/mol.
Problem 3.31: Given values ∆H1, ∆H2, ∆egH for several elements, classify reactivity types and compounds they form (MX2, MX, etc.).
Problem 3.34: Identify an incorrect statement about the modern periodic table.
Problem 3.39: Order of non-metallic character among elements F, Cl, O, N; the correct order is typically F > Cl > O > N in terms of electronegativity and oxidizing capability.
Problem 3.40: Order of chemical reactivity in terms of oxidizing strength among F, Cl, O, N: F > Cl > O > N (most to least oxidizing).