Notes on Classification of Elements and Periodicity in Properties
3.1 Why Do We Need to Classify Elements?
Elements are the basic units of all matter. Historically, as the number of known elements grew from 31 (in 1800) to 63 (by 1865) and now to 114, it became clear that organizing them systematically would ease study and enable predictions about undiscovered elements. A classification helps rationalize known facts and predict new ones.
The Periodic Table provides a concise organization of chemistry, showing trends and grouping elements into families rather than treating them as a random collection.
Objectives of this unit include understanding how grouping by properties led to the Periodic Table, the Modern Periodic Law, the significance of atomic number and electronic configuration, IUPAC nomenclature for Z>100, classification into s, p, d, f blocks and their main characteristics, periodic trends in physical/chemical properties, and the relation of ionization enthalpy to metallic character. It also encourages proper scientific vocabulary (atomic/ionic radii, ionization enthalpy, electron gain enthalpy, electronegativity, valence).
3.2 Genesis of Periodic Classification
The idea of trends in properties emerged with several scientists:
Johann Dobereiner (early 1800s) proposed Triads: groups of three elements where the middle element’s properties lie between the other two and its atomic weight is about halfway between the others. This led to the concept of periodic recurrence of properties (Table 3.1).
John Alexander Newlands (1865) proposed the Law of Octaves: in increasing order of atomic weights, every eighth element resembled the first element (like musical octaves). However, it worked well only up to calcium and was not universally accepted.
A French geologist, A. E. B. de Chancourtois (1862), proposed a cylindrical arrangement displaying repeating properties, a precursor to a periodic arrangement.
Lothar Meyer, independently around 1869, plotted physical properties (atomic volume, melting/boiling points) against atomic weight and observed periodic patterns; his table resembled the modern periodic table, but his work was not published before Mendeleev.
Dmitri Mendeleev (1834–1907) is widely credited with publishing the Periodic Law for the first time. He arranged elements by increasing atomic weights, placing elements with similar properties in the same vertical groups, and left gaps for undiscovered elements, predicting their properties (e.g., Eka-aluminium = Gallium, Eka-silicon = Germanium). This bold predictive power contributed to his fame.
The concept of periodicity was strengthened by Mendeleev’s reliance on chemical formulas and properties of compounds, not just atomic weights.
The progression culminated in the modern view that periodicity arises from electronic structure, not just atomic weights.
Key historical milestones (names and ideas):
Dobereiner's Triads: middle element ~ average of the other two; properties in-between.
Newlands’ Octaves: every eighth element similar in properties (up to Ca).
Chancourtois’ cylinder/tableau: early attempt at a 3D arrangement.
Meyer’s periodic tables based on physical properties vs. atomic weight.
Mendeleev’s periodic law and table, with predictive gaps.
3.3 Modern Periodic Law and the Present Form of the Periodic Table
A major shift occurred with the discovery of the atomic nucleus and sub-atomic structure. In 1913, Henry Moseley showed that a plot of X-ray frequency vs. atomic number produced a straight line, revealing that the atomic number (Z), i.e., the nuclear charge, is the fundamental property that governs the periodic recurrence of properties, not the atomic weight.
Modern Periodic Law: The physical and chemical properties of the elements are periodic functions of their atomic numbers:
The periodic table now reflects electronic configurations. The periodicity is essentially a consequence of the periodic variation in electronic configurations, which determine properties.
The modern long-form periodic table is the most convenient: horizontal rows are periods and vertical columns are groups. Elements with similar outer electron configurations occupy the same group.
IUPAC has standardized group numbering from 1 to 18 (replacing the older IA–VIIA, VIII, IB–VIIB, 0). There are seven periods in total. The periods have the following maximum possible element counts based on electron capacity: 2 (1st), 8 (2nd), 8 (3rd), 18 (4th), 18 (5th), 32 (6th), and 32 (7th theoretical maximum; the seventh is incomplete).
The seventh period would include most of the transactinide elements and is incomplete; the sixth period includes the lanthanoids, and the seventh period includes the actinoids, with lanthanoid and actinoid series placed in separate bottom panels in the long form to preserve clear periodic trends.
A modern form of the table emphasizes outer electronic configurations; the noble gases occupy a closed-shell configuration and are very unreactive.
Historical context about the boundaries: discovery of short-lived elements led to the extension of the table and recognition that internal structure (electronic configuration) drives periodicity.
3.4 Nomenclature of Elements with Atomic Numbers > 100
The naming of new elements, especially with very high atomic numbers, is a source of controversy because such elements are highly unstable and produced only in minute quantities.
Until a discovery is confirmed and names are officially recognized, IUPAC uses a systematic nomenclature based on the atomic number, using numerical roots for digits 0–9 and adding the suffix -ium. Roots (Table 3.4) map digits to syllables: 0 nil, 1 un, 2 bi, 3 tri, 4 quad, 5 pent, 6 hex, 7 sept, 8 oct, 9 enn. The stems are combined in the order of the digits, followed by -ium (e.g., Ununtrium for Z=113, etc.).
The IUPAC temporary name and symbol are three-letter, then later replaced by a permanent name and symbol once a consensus is reached (often referencing a country, state, or notable scientist).
As of the period of the text, elements up to Z = 118 have been discovered and officially named.
Example: Problem 3.1 asks for the IUPAC name and symbol for Z = 120. Using roots 1 (un), 2 (bi), and 0 (nil) gives Unbinilium with symbol Ubn.
Tables (Table 3.4 and Table 3.5) illustrate the nomenclature with the sequence of roots and the corresponding temporary names and symbols, and the eventual permanent names and symbols for those elements.
3.5 Electronic Configurations of Elements and the Periodic Table
An electron in an atom is characterized by a set of quantum numbers; the principal quantum number n defines the main energy level (shell). The distribution of electrons into subshells (s, p, d, f) is the electronic configuration.
The location of an element in the modern periodic table reflects the quantum numbers of the last orbital filled. The long form (current standard) shows a simple relation between periods and the filling of shells.
Periods correspond to the principal quantum number n of the outermost shell:
1st period (n = 1): 2 elements (H, He) with configuration 1s$^1$ and 1s$^2$.
2nd period (n = 2): 8 elements (Li to Ne); filling 2s and 2p orbitals.
3rd period (n = 3): 8 elements (Na to Ar).
4th period (n = 4): 18 elements (K to Kr); 3d orbitals begin to fill (Sc to Zn) before 4p is completed.
5th period (n = 5): 18 elements (Rb to Xe); 4d transition series starts (Y to Ag) and 5p fills to Xe.
6th period (n = 6): 32 elements; filling order is 6s, 4f, 5d, 6p; 4f orbitals begin filling at Ce (Z = 58) and end at Lu (Z = 71) (lanthanides) forming the 4f inner-transition series.
7th period (n = 7): similar to the 6th; includes 7s, 5f, 6d, 7p, and many transactinide elements; ends near Og (Z = 118) in this period.
The 4f and 5f shells form the lanthanoid and actinoid series, which are placed in separate panels in the long form to maintain a clear layout.
The section includes several worked examples and Problems (3.1–3.6, etc.) that illustrate the use of electronic configurations to justify placement in a given group or period.
Examples mentioned include the hypothetical placements of Ga and Ge (in early versions) and the known modern placements according to Z. The section also notes that the 7th period would include a theoretical maximum of 32 elements, similar to the 6th period, though it is incomplete in practice.
3.6 Electronic Configurations and the Types of Elements: s-, p-, d-, f- BLOCKS
The Aufbau principle and the filling of subshells give rise to four blocks in the periodic table, corresponding to the type of orbital being filled:
s-block: elements with outer electronic configuration ns$^1$ or ns$^2$ (Group 1 and Group 2).
p-block: elements with outer configuration ns$^2$np$^1$ to ns$^2$np$^6$ (Groups 13–18). The p-block and s-block together form the Representative (Main Group) Elements.
d-block: Transition elements (Groups 3–12); characterized by filling inner d orbitals; general outer configuration is (n-1)d$^{1-10}$ns$^{0-2}$ (except Pd: 4d$^{10}$5s$^0$).
f-block: Inner-transition elements (Lanthanoids Ce–Lu; Actinoids Th–Lr) with outer configuration (n-2)f$^1$–f$^{14}$ (n-1)d$^0$–d$^1$ ns$^2$; these elements are metallic and typically radioactive.
He is an exception: though placed in the s-block, He is placed with the noble gases in the p-block due to its full valence shell (1s$^2$).
Hydrogen is a special case: it has one 1s electron and can behave like an alkali metal (Group 1) or gain an electron to resemble a halogen (Group 17). For practical purposes, H is placed separately at the top of the periodic table.
The text also discusses the four major classifications: Metals, Non-metals, and Metalloids, and situates them within the blocks.
3.6.1 The s-Block Elements
Includes Group 1 (alkali metals) with ns$^1$ configuration and Group 2 (alkaline earth metals) with ns$^2$ configuration.
They are highly reactive metals with low ionization enthalpies; they lose their outer electrons readily to form 1+ (alkali) or 2+ (alkaline earth) ions. Reactivity increases down the group; they are not found free in nature due to high reactivity.
Compounds of s-block elements are predominantly ionic, with the exception of Li and Be which show more covalent character.
Representative examples: Li, Na, K, Rb, Cs, Fr (Group 1) and Be, Mg, Ca, Sr, Ba, Ra (Group 2).
3.6.2 The p-Block Elements
The p-block includes Groups 13–18 and contains the Representative Elements (Main Group Elements).
Outer configurations vary from ns$^2$np$^1$ to ns$^2$np$^6$ in each period. End of each period features a noble gas with ns$^2$np$^6$.
Noble gases are highly unreactive due to complete valence shells.
Two chemically important non-metals precede the noble gases: the Halogens (Group 17) and the Chalcogens (Group 16). They have highly negative electron gain enthalpies and readily gain electrons to reach noble gas configurations.
Across a period, metallic character decreases to the right; non-metallic character increases to the right. The line separating metals from non-metals is not sharp; metalloids (Si, Ge, As, Sb, Te) lie along the diagonal boundary and show mixed properties.
He is an exception in the s-block, and H is a special case in the p-block; these exceptions are discussed in 3.6.2.
3.6.3 The d-Block Elements (Transition Elements)
The d-block comprises Groups 3–12 and includes transition metals.
Characterized by filling of inner d orbitals, outer configuration typically (n-1)d$^{1-10}$ns$^{0-2}$ (except Pd: 4d$^{10}$5s$^0$).
Properties: metals, colored ions, variable oxidation states, paramagnetism, and frequent catalytic behavior.
Zn, Cd, and Hg (with (n-1)d$^{10}$ns$^2$) are exceptions and do not show typical transition-element behavior.
Transition metals act as a bridge between highly reactive s-block metals and the p-block elements; often exhibit catalytic properties.
3.6.4 The f-Block Elements (Inner-Transition Elements)
The bottom two rows of the table: Lanthanoids (Ce–Lu, Z = 58–71) and Actinoids (Th–Lr, Z = 90–103).
Outer configuration: (n-2)f$^1$–f$^{14}$ (n-1)d$^0$–d$^1$ ns$^2$.
They are all metals and have similar chemistry within each series.
Actinoids are generally radioactive and many have been produced only in trace quantities; their chemistry is not fully explored.
Elements after uranium are called Transuranium Elements.
3.6.5 Metals, Non-metals and Metalloids
A broad classification shown in Fig. 3.3: Metals (left side, >78% of known elements), Non-metals (top-right), and Metalloids (borderline region along the diagonal line between metals and non-metals).
Properties:
Metals: typically solid at room temperature (mercury is an exception), high melting/boiling points, good conductors, malleable, ductile.
Non-metals: often solids or gases at room temperature, have low melting/boiling points, poor conductors, brittle in solid form.
Metalloids: exhibit mixed properties with a zig-zag boundary between metals and non-metals.
The metallic character increases down a group and decreases across a period; non-metallic character behaves oppositely.
3.7 Periodic Trends in Properties of Elements
There are many observable trends in physical and chemical properties as we move across a period or down a group. Here we discuss key trends in terms of electrons and energy levels.
3.7.1 Trends in Physical Properties
(a) Atomic Radius
Atomic size is not directly measurable; a practical estimate is the atomic radius defined as half the distance between two bound atoms in a molecule (covalent radius) or half the distance between two adjacent atoms in a metallic crystal (metallic radius).
Practical examples: Covalent radius of Cl is half the Cl–Cl bond distance in Cl2: bond distance ≈ 198 pm, so radius ≈ 99 pm. Metallic radius for Cu is half the interatomic distance in solid Cu: Cu–Cu distance ≈ 256 pm, so radius ≈ 128 pm.
Two clear trends:
Across a Period: radius generally decreases from left to right. Reason: as Z increases, outer electrons are in the same principal energy level (same n) while the effective nuclear charge increases, pulling electrons closer.
Down a Group: radius increases down the group. Reason: larger principal quantum number n for outer electrons and shielding by inner electrons reduces effective nuclear attraction.
Note: Noble gases radii are typically discussed as van der Waals radii; monoatomic gases differ from covalent radii.
Table 3.6 (selected values): Across Period II (Period 2): Li 152, Be 111, B 88, C 77, N 74, O 66, F 64 (pm). Across Period III (Period 3): Na 186, Mg 160, Al 143, Si 117, P 110, S 104, Cl 99 (pm). Down a Group I (Group 1): Li 152, Na 186, K 231, Rb 244, Cs 262, At 140 (pm).
(b) Ionic Radius
Cations are smaller than their parent atoms; anions are larger due to electron removal/addition effects and shielding.
Isoelectronic species (same electron count) can have different radii; larger nuclear charge leads to smaller radius.
Example: F$^-$ (136 pm) vs F (64 pm); Na$^+$ (95 pm).
Concepts of isoelectronic species and trends explained with these ideas.
Fig. 3.4 illustrates the variation of R with Z across periods and down groups.
(c) Ionization Enthalpy (First Ionization Enthalpy, ΔiH)
The energy required to remove an electron from a neutral gaseous atom: with as the enthalpy change. Units: kJ mol$^{-1}$. The second ionization enthalpy corresponds to removing the second electron: . More generally, successive ionization enthalpies increase.
The first ionization enthalpy generally increases across a period and decreases down a group.
Maxima at noble gases (stable closed shells); minima at alkali metals (low ionization enthalpy).
Trends explained by two factors:
(i) Attraction of electrons to the nucleus (effective nuclear charge, Z_eff).
(ii) Electron shielding by inner cores.
Key anomaly discussions include Be vs B (2s vs 2p penetration and shielding) and O vs N (Hund’s rule and electron-electron repulsion in 2p orbitals).
Figure references: Fig. 3.5 (ΔiH vs Z for Z = 1–60); Fig. 3.6(a) ΔiH for second-period elements; Fig. 3.6(b) ΔiH for alkali metals.
Problem 3.6: The first ionization enthalpy values for Na, Mg, Si are 496, 737, 786 kJ/mol, respectively. The Al value would be closer to which: 575 or 760? Answer: ~575 kJ/mol, due to 3p shielding by 3s electrons.
(d) Electron Gain Enthalpy (Electron Affinity, ΔegH)
The enthalpy change when an electron is added to a neutral gaseous atom:
Electron affinity is exothermic (negative ΔegH) for many nonmetals (especially halogens) because they gain electrons to reach noble-gas configuration. Noble gases have positive electron affinities because adding an electron would enter a higher-energy, unstable shell.
Across a period, electron gain enthalpy generally becomes more negative; down a group, it generally becomes less negative due to increasing atomic size and weaker attraction for added electrons.
Table 3.7 (selected): For Group 1: H = –60 to –73 range; Li = –60; Na = –53; K = –48; Rb = –47; Cs = –46. For Group 17: F = –328; Cl = –349; Br = –325; I = –295; At = –270. For Group 16: O = –141; S = –200; Se = –195; Te = –190; Po = –174. Ne, Ar, Xe, Rn show positive values (e.g., Ne +116, Ar +96, Xe +77, Rn +68).
The trend is not strictly systematic across all elements but is linked to effective nuclear charge and orbital occupancy.
(e) Electronegativity
Electronegativity is a qualitative (not directly measurable) measure of an atom’s ability to attract shared electrons in a chemical bond.
Pauling scale (most widely used): F is assigned 4.0; across a period, electronegativity generally increases from left to right; down a group, it decreases.
Periodic trends relate electronegativity to atomic radii and oxidation tendency: smaller atoms (higher Z_eff) attract electrons more strongly; non-metals tend to attract electrons more than metals.
Values (selected):
Period II: Li 1.0, Be 1.5, B 2.0, C 2.5, N 3.0, O 3.5, F 4.0
Period III: Na 0.9, Mg 1.2, Al 1.5, Si 1.8, P 2.1, S 2.5, Cl 3.0
Group I: Li 1.0, Na 0.9, K 0.8, Rb 0.8, Cs 0.7
Group 17: F 4.0, Cl 3.0, Br 2.8, I 2.5, At 2.2
The electronegativity trend mirrors the metallic vs. non-metallic character: increases across a period (more non-metallic) and decreases down a group (more metallic).
3.7.2 Periodic Trends in Chemical Properties
Valence (oxidation) states: The valence of representative elements often equals the number of electrons in the outermost shell, or eight minus that number, yielding common valence states and oxo/hydride patterns. E.g., SiH$_4$ (valence 4), NaH (valence 1).
The oxidation state concept extends to the oxide and hydride formation (see Table 3.9 for periodic trends in valence across groups 1, 2, 13–17).
Anomalous second-period properties: the first element of Groups 1 (Li) and 2 (Be) and Groups 13–17 (B–F) show deviations from the rest of their groups due to small size, high charge-to-radius ratio, and high electronegativity; e.g., Boron forms covalent compounds like BF$_3$ with a covalency up to 4, whereas heavier members can expand their valence beyond four.
Diagonal relationship: similarities in properties between certain pairs of elements diagonally adjacent in the table, e.g., Mg and Al share some chemical behavior despite being in different groups.
Problem 3.8: Predict formulas from periodic trends (Si + Br → SiBr$4$; Al + S → Al$2$S$_3$).
Problems 3.9 and others illustrate the difference between oxidation state and covalency (e.g., in [AlCl(H$2$O)$5$]$^{2+}$, Al has +3 oxidation state but covalency 6).
3.7.3 Periodic Trends and Chemical Reactivity
Reactivity shows a maximum at the two extremes of a period (alkali metals at left and halogens at right) and a minimum in the middle. This correlates with the propensity to lose or gain electrons:
Left edge: high metallic character and low ionization enthalpy; tendency to lose electrons (reducing agents).
Right edge: high non-metallic character and tendency to gain electrons (oxidizing agents).
Within a group of metallic elements, reactivity generally increases down the group; within a group of non-metals (halogens), reactivity decreases down the group.
Oxidation with oxygen varies: leftmost elements form basic oxides (e.g., Na$2$O), rightmost form acidic oxides (e.g., Cl$2$O$7$). Middle oxides can be amphoteric (e.g., Al$2$O$_3$) or neutral (e.g., CO).
The trend in the transition metals is more nuanced due to d-orbital participation; radii changes are smaller and ionization enthalpies are intermediate between s- and p-blocks.
Problem 3.10 asks to show Na$2$O is a basic oxide and Cl$2$O$7$ is an acidic oxide by reaction with water: Na$2$O + H$2$O → 2 NaOH; Cl$2$O$7$ + H$2$O → 2 HClO$_4$.
3.8 Summary of Periodic Trends and Concepts
The Periodic Law and Periodic Table originated from the quest to organize elements by recurring properties and later by electronic structure.
The modern Periodic Table arranges elements by increasing atomic number in seven periods and 18 groups, with groups indicating similar valence electron configurations.
Four block types (s, p, d, f) correspond to the type of orbital filled in the outer shell (with He and H as special cases).
The table shows broad classes: Metals, Non-metals, and Metalloids, with metallic character increasing down a group and decreasing across a period; non-metallic character showing opposite trends.
Periodic trends include: atomic/ionic radii, ionization enthalpy, electron gain enthalpy, electronegativity, and valence behavior. The relationships among these properties explain general conductivity, reactivity, and bonding tendencies.
The chemistry of elements on the left (alkali/alkaline earth metals) tends toward oxidation via electron loss, whereas the right-hand side (halogens, chalcogens) tends toward electron gain.
The diagonal relationship explains certain similarities between certain pairs of elements across groups (e.g., Mg and Al).
Problems and Solutions (Selected Examples)
Problem 3.1
Question: What would be the IUPAC name and symbol for the element with atomic number 120?
Solution: Using Table 3.4 roots for 1, 2, 0 as un, bi, nil respectively, the temporary name would be Unbinilium with symbol Ubn. (Unbinilium, Ubn)
Problem 3.2
Question: 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 available orbitals 5s, 4d, and 5p increases is 5s < 4d < 5p. The total number of orbitals available is 9, so the maximum number of electrons that can be accommodated is 18; hence 18 elements in the 5th period.
Problem 3.3
Question: Elements Z = 117 and 120 have not yet been discovered. In which family/group would you place these elements and what would their electronic configurations be?
Solution: Z = 117 would belong to the halogen family (Group 17) with electronic configuration
(shown as [Rn] 5f$^{14}$6d$^{10}$7s$^2$7p$^5$). Z = 120 would be placed in Group 2 (alkaline earth metals) with electronic configuration
Problem 3.4
Question: Justify the presence of 32 elements in the 6th period (and the theoretical maximum for the 7th period).
Solution: For the 6th period (n = 6), the allowed orbitals are 6s, 4f, 5d, 6p. The total number of orbitals is 2 (6s) + 14 (4f) + 10 (5d) + 6 (6p) = 32. This yields 32 elements in the 6th period. The 7th period would similarly involve 7s, 5f, 6d, 7p orbitals, with a theoretical maximum of 32 elements; in practice, the period is incomplete due to the limited experimental discovery of heavier transactinide elements.
Problem 3.5
Question: In terms of period and group where would you locate the element with Z = 114?
Solution: Z = 114 lies in the 7th period and the p-block; more specifically, it belongs to Group 14 (carbon family) in the long-form periodic table. (Note: this is delivered in the chapter’s problem set with the expected placement for Z = 114.)
Problem 3.6
Question: The first ionization enthalpies (∆iH) of Na, Mg, Si are 496, 737, 786 kJ mol$^{-1}$ respectively. Predict whether the first ∆iH for Al will be closer to 575 or 760 kJ mol$^{-1}$? Justify.
Solution: It will be closer to 575 kJ mol$^{-1}$. Across the same period, shielding by the inner electrons is not significantly increased as you move from Na to Mg to Si; however, when moving to Al from Mg, the 3p electron experiences greater shielding by the 3s electrons, reducing the effective nuclear charge felt by the valence electron, causing ∆iH(Al) to be lower than Mg. Thus, around 575 kJ mol$^{-1}$ is more reasonable than 760 kJ mol$^{-1}$.
Problem 3.7
Question: Which of the following will have the most negative electron gain enthalpy and which the least: P, S, Cl, F? Explain.
Solution: Electron gain enthalpy generally becomes more negative across a period from left to right. Among the listed elements, Cl has the most negative electron gain enthalpy (most negative, strongest tendency to gain an electron). P has the least negative electron gain enthalpy (least tendency to gain an electron among the four). This is consistent with the typical trend across a period and the fact that halogens achieve noble gas configurations most readily by gaining one electron.
Problem 3.8
Question: Using the Periodic Table, predict the formulas of compounds formed by the following pairs of elements:
(a) silicon and bromine
(b) aluminium and sulphurSolution:
(a) Si is a Group 14 element with valence 4; Br is a Group 17 element with valence 1. The formula would be SiBr$4$. (b) Al is Group 13 with valence 3; S is Group 16 with valence 2. The formula would be Al$2$S$_3$.
Problem 3.9
Question: Are the oxidation state and covalency of Al in [AlCl(H$2$O)$5$]$^{2+}$ the same?
Solution: No. The oxidation state of Al is +3, but the covalency is 6, reflecting complex coordination by water ligands and chloride in this hexacoordinate complex.
Problem 3.10
Question: Show by a chemical reaction with water that Na$2$O is a basic oxide and Cl$2$O$_7$ is an acidic oxide.
Solution: Na$2$O + H$2$O → 2 NaOH (basic oxide). Cl$2$O$7$ + H$2$O → 2 HClO$4$ (acidic oxide).
Appendix: Key Tables and Formulas (Selected)
Atomic radii (pm): Across Period II: Li 152, Be 111, B 88, C 77, N 74, O 66, F 64. Across Period III: Na 186, Mg 160, Al 143, Si 117, P 110, S 104, Cl 99. Down Group I: Li 152, Na 186, K 231, Rb 244, Cs 262, At 140.
Electron gain enthalpies ΔegH (kJ/mol): H –60 to –73, Li –60, Na –53, K –48, Rb –47, Cs –46; O –141, S –200, Se –195, Te –190, Po –174, F –328, Cl –349, Br –325, I –295, At –270; Ne +116, Ar +96, Xe +77, Rn +68.
Electronegativity (Pauling scale): Period II (Li 1.0, Be 1.5, B 2.0, C 2.5, N 3.0, O 3.5, F 4.0); Period III (Na 0.9, Mg 1.2, Al 1.5, Si 1.8, P 2.1, S 2.5, Cl 3.0); Group I (Li 1.0, Na 0.9, K 0.8, Rb 0.8, Cs 0.7); Group 17 (F 4.0, Cl 3.0, Br 2.8, I 2.5, At 2.2).
Problem 3.15 (Hydrogen energy reference): The energy of the electron in the ground state of hydrogen is The ionization enthalpy of atomic hydrogen can be derived via the mole concept.
The IUPAC nomenclature roots (Table 3.4) for digit 0–9: 0 = nil, 1 = un, 2 = bi, 3 = tri, 4 = quad, 5 = pent, 6 = hex, 7 = sept, 8 = oct, 9 = enn;Elements above 100 are named with these roots: Unnilunium, Ununbium, etc., followed by -ium (e.g., Ununquadium for Z = 114; later renamed as Flerovium, etc.).
Quick Reference: Key Concepts
Modern Periodic Law: The properties of elements are periodic functions of their atomic numbers, driven by electronic configurations.
Periods and Groups: Periods correspond to shells being filled (n), Groups to similar outer electron configurations.
Four blocks: s-block (ns$^1$, ns$^2$), p-block (ns$^2$np$^{1-6}$), d-block (inner transition metals), f-block (lanthanoids and actinoids).
Periodic trends: Atomic and ionic radii, ionization enthalpy, electron gain enthalpy, and electronegativity show systematic variation across periods and down groups, with characteristic anomalies (Be/B, O/N, Li/Be, etc.).
Reactivity pattern: High reactivity at extremes of a period (metallic left and non-metallic right); metallic character increases down a group, non-metallic character decreases down a group.
Oxidation states: Valence approximates the outer-shell electrons; oxidation states help predict formulas of compounds (e.g., NaH, LiH, Al$2$O$3$, etc.).
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