Classification of Elements - The Periodic Table
1. What This Chapter Covers
A medical shop holds a vast number of medicines, and the shopkeeper cannot possibly remember every name. Yet when you ask for one, he hands it over without difficulty. Step into a super bazaar and the same thing is true — you find the grocery you need because the goods are arranged in a particular order.
The chapter opens with those two observations and draws one conclusion from them: for any system involving several things, a particular order of arrangement is essential.
Chemistry faced exactly that problem. Robert Boyle (1661) defined an element as any substance that cannot be decomposed into a simpler substance by a physical or chemical change. About thirteen elements were known at his time.
The count then grew fast:
| By when | Elements known |
|---|---|
| Boyle, 1661 | about 13 |
| End of the eighteenth century, Lavoisier's time | another 11 discovered |
| 1865 | about 63 |
| 1940 | 91 from natural sources plus 17 synthetic |
| Today, including synthetic elements | more than 118 |
As the number grew, it became impossible to keep in memory the chemistry of individual elements and their compounds. The older division into metals and non-metals had clear limits — aluminium shows some metallic and some non-metallic properties, and elements of that type had to be called metalloids or semi-metals. A better classification was needed.
In the beginning of the eighteenth century Joseph Louis Proust stated that the hydrogen atom is the building material, and the atoms of all other elements are simply combinations of some number of hydrogen atoms. At his time all atomic weights were given as whole numbers with hydrogen taken as one, which is what made the idea look reasonable.
The chapter is allotted 10 periods in September and runs from textbook page 123 to page 149.
2. Dobereiner's Law of Triads
A German chemist, Johann Wolfgang Dobereiner (1829), noticed groups of three elements — triads — whose members had similar chemical properties. He tried to connect those properties to atomic weights.
His statement, the law of triads, is that when elements with similar properties are taken three at a time and arranged in ascending order of their atomic weights, the atomic weight of the middle element is the average of the atomic weights of the first and third.
Activity 1 gives five such groups to check:
| Group | First | Middle | Third |
|---|---|---|---|
| A | Lithium (Li) 7.0 | Sodium (Na) 23.0 | Potassium (K) 39.0 |
| B | Calcium (Ca) 40.0 | Strontium (Sr) 88 | Barium (Ba) 137.0 |
| C | Chlorine (Cl) 35.5 | Bromine (Br) 80.0 | Iodine (I) 127.0 |
| D | Sulphur (S) 32.0 | Selenium (Se) 78.0 | Tellurium (Te) 125.0 |
| E | Manganese (Mn) 55.0 | Chromium (Cr) 52.0 | Iron (Fe) 56.0 |
The worked case the book prints is row A: (7.0 + 39.0) / 2 = 23.0, which is exactly the atomic weight of sodium. The activity asks you to do the same arithmetic for the remaining rows and compare each result with the middle element.
Dobereiner's attempt mattered less for its arithmetic than for the clue it gave — that atomic weights could be correlated with properties. It made chemists start looking at elements in terms of groups with similar chemical and physical behaviour, and that eventually led to the modern table.
Limitations
- Not all the known elements of the time could be arranged as triads.
- The law failed for very low mass and for very high mass elements. For F, Cl and Br the atomic weight of Cl is not the arithmetic mean of F and Br.
- As techniques for measuring atomic masses improved, the law stopped being strictly valid.
The chapter's Think and discuss here asks what relation Dobereiner wanted to establish, and then sets a problem worth doing: the densities of calcium and barium are 1.55 and 3.51 g cm⁻³, so what does the law of triads predict for strontium?
3. Newlands' Law of Octaves
John Newlands, a British chemist, found in 1865 that when elements were arranged in ascending order of atomic weight they appeared to fall into seven groups, each containing elements of similar chemical properties.
The law of octaves states that when elements are arranged in ascending order of their atomic weights, their properties repeat at regular intervals — every eighth element, starting from a given element, resembles that element in its properties. Elements with similar chemical properties are placed along a horizontal row.
Table 1 in the book is Newlands' table of 1866, numbering the elements from H = 1 down to Th = 56. Newlands was the first to assign numbers to the elements. His work was accepted neither by his seniors nor by the Journal of the Chemical Society, which refused to publish it.
Reading the table as he intended: start at hydrogen, move down to oxygen, then start again at the top — the eighth element from hydrogen is fluorine, the next eighth is chlorine, and the properties of hydrogen, fluorine and chlorine are similar. Start at lithium and the eighth is sodium, the next potassium; all three are reactive metals.
The table had real problems:
- Two elements were fitted into the same slot, for example cobalt and nickel.
- Elements totally dissimilar in properties were put in the same group. Newlands placed Co, Ni, Pd, Pt and Ir in the same row as the halogens F, Cl, Br and I.
- The table was restricted to 56 elements and left no room for new ones, so later discoveries could not be fitted in according to their properties.
The reason he forced the pattern is in the name. Newlands was trying to link the periodicity of chemical properties to the periodicity found in music, where any note in a key is separated from its octave by an interval of seven notes.
The book's Do you know? box makes the parallel explicit: the Indian scale has seven notes — sa, re, ga, ma, pa, da, ni — and the western one do, re, mi, fa, so, la, ti, with every eighth note similar to the first and beginning the next scale. Chasing that analogy made him sometimes ignore the similarities.
4. Mendeleeff's Periodic Table
Mendeleeff arranged the elements known at his time in increasing order of atomic weight, in a chart divided into 8 vertical columns called groups, each group split into A and B sub-groups. Elements in a group have similar properties; elements along a period differ.
The evidence he used was the formula of the oxide. Elements of the first column react with oxygen to give compounds of the general formula R₂O — Li₂O, Na₂O, K₂O. Elements of the second column give RO — BeO, MgO, CaO. Mendeleeff explained the similarity within a group in terms of common valency.
The periodic law
From these observations came the law that carries his name:
The physical and chemical properties of the elements are periodic functions of their atomic weights.
Salient features and achievements
1. Groups and sub-groups. Eight vertical columns, numbered with Roman numerals I to VIII, each divided into sub-groups A and B. Elements within a sub-group resemble one another more closely — sub-group IA, the alkali metals Li, Na, K, Rb and Cs, are a good example.
2. Periods. The horizontal rows, denoted 1 to 7. Elements in a period differ in properties from one another. A period is the entire range of elements after which properties repeat.
3. Predicting missing elements. He left blank spaces where he believed elements were still undiscovered, and predicted their properties in advance from the table alone. He named them by adding the prefix eka — Sanskrit for the numeral one — to the element immediately above the empty space. Eka-boron, eka-aluminium and eka-silicon turned out to be scandium, gallium and germanium.
Table-3 sets his predictions beside the measured values:
| # | Property | Eka-Aluminium (Ea) | Eka-Silicon (Es) | Gallium (1875) | Germanium (1886) |
|---|---|---|---|---|---|
| 1 | Atomic weight | 68 | 72 | 69.72 | 72.59 |
| 2 | Density | 5.9 | 5.5 | 5.94 | 5.47 |
| 3 | Formula of oxide | Ea₂O₃ | EsO₂ | Ga₂O₃ | GeO₂ |
| 4 | Formula of chloride | EaCl₃ | EsCl₄ | GaCl₃ | GeCl₄ |
The book adds one more prediction in a Do you know? box. Asked about the melting point of eka-aluminium, Mendeleeff said "If I hold it in my hand, it will melt" — gallium melts at 30.2 °C and body temperature is 37 °C.
4. Correcting atomic weights. The table itself corrected the atomic masses of beryllium, indium and gold. Beryllium was then given the atomic weight 13.5, because atomic weight = equivalent weight × valency, its equivalent weight was measured as 4.5 and its valency was thought to be 3. That weight put it in the wrong group.
Mendeleeff argued the valency must be 2, giving 4.5 × 2 = 9, which places it in the second group beside Mg and Ca — where its properties actually belong.
5. Anomalous series. In places an element of higher atomic weight had to be placed before one of lower weight, as with Te (127.6 u) before I (126.9 u). Mendeleeff accepted these minor inversions because they put the elements in the correct groups.
It was this extraordinary thinking that made chemists accept the periodic table and recognise Mendeleeff, more than anyone else, as the originator of the periodic law. He did all of it before the electron had even been discovered, and the 101st element was named mendelevium in his honour.
5. Limitations of Mendeleeff's Table
1. Anomalous pairs of elements. Certain elements of higher atomic weight precede those with lower atomic weight — tellurium, 127.6, before iodine, 126.9.
2. Dissimilar elements placed together. Elements with unlike properties sat in the same group as sub-groups A and B. The alkali metals Li, Na and K of IA have little resemblance to the coinage metals Cu, Ag and Au of IB. Cl is in VIIA and Mn in VIIB — but chlorine is a non-metal and manganese a metal.
The Think and discuss boxes in this stretch are worth answering in writing: why Mendeleeff had to leave blank spaces, what Ea₂O₃ and EsO₂ mean, and whether hydrogen — a diatomic gas — really belongs in the first group with the solid alkali metals.
6. Moseley and the Modern Periodic Law
H.J. Moseley (1913) found that each element emits a characteristic pattern of X-rays when bombarded with high-energy electrons. By analysing those patterns he could calculate the number of positive charges in the atom of each element.
The number of positive charges — protons — in the atom of an element is its atomic number. Moseley realised that the atomic number Z is a more fundamental characteristic of an element than its atomic weight.
Arranging elements by increasing atomic number eliminated the anomalous series at a stroke. Tellurium has a greater atomic weight than iodine, but its atomic number is one unit less, so the order comes out right by itself.
That forced the law to change:
Mendeleeff's periodic law: the properties of elements are periodic functions of their atomic weights.
Modern periodic law: the properties of the elements are periodic functions of their atomic numbers.
There is one further refinement. The atomic number gives both the protons in the nucleus and the electrons in the neutral atom — and the physical and chemical properties depend not on the protons but on the number of electrons and their arrangement. So the law can also be stated as:
The physical and chemical properties of elements are the periodic functions of the electronic configurations of their atoms.
The table built on this law is the long form of the periodic table, an extension of Mendeleeff's original short form.
7. Where an Element Sits in the Modern Table
The modern periodic table has eighteen vertical columns called groups and seven horizontal rows called periods.
The classification rests on electron arrangements that repeat periodically. Elements with similar outer shell (valence shell) configurations fall in the same column; going down a group, the principal quantum number increases.
From Chapter 6 you know the capacities of the sub-shells: s has one orbital and holds 2 electrons, p has 3 orbitals and holds 6, d has 5 orbitals and holds 10, and f has 7 orbitals and holds 14.
Elements are classified as s, p, d or f block according to which sub-shell the differentiating electron — the last electron to enter — goes into:
| Element | Configuration (last electron underlined in the book) | Block |
|---|---|---|
| ₁₁Na | 1s² 2s² 2p⁶ 3s¹ | s |
| ₁₃Al | 1s² 2s² 2p⁶ 3s² 3p¹ | p |
| ₂₁Sc | 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹ | d |
| ₅₈Ce | 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 5d¹ 4f² | f |
Table-4 in the book lays the same four configurations out as a grid of shells and sub-shells, and is a useful cross-check on your own filling.
Groups
The eighteen groups were traditionally written with Roman numerals I to VIII plus the letters A and B. The latest IUPAC recommendation is Arabic numerals 1 through 18 with no A or B. The book uses the new numbering with the old one in parentheses — Group 2 (IIA), Group 16 (VIA).
A group is also called an element family or chemical family. Group 1 (IA), from Li to Fr, has the outer configuration ns¹ and is the alkali metal family. Activity 2 asks you to complete Table-5 for the main-group families with their range of elements, valence electrons, valence shell configuration and valency:
| Group | Family |
|---|---|
| 1 (IA) | Alkali metal family |
| 2 (IIA) | Alkaline earth metal family |
| 13 (IIIA) | Boron family |
| 14 (IVA) | Carbon family |
| 15 (VA) | Nitrogen family |
| 16 (VIA) | Oxygen or chalcogen family |
| 17 (VIIA) | Halogen family |
| 18 (VIIIA) | Noble gas family |
The names come from the chemistry, as the Do you know? box explains. Aliquili means plant ash, and Na and K were obtained from plant ash. Chalcogenous means ore product, because group 16 elements form ores with metals. Halos is sea salt and genus is produced, because most group 17 elements come from nature as sea salt. The noble gases are named for being chemically least active, and their outer shell configurations are the basis of the octet rule.
Periods
Seven periods, numbered 1 to 7. Two rules decide them:
- The number of main shells in the atom decides the period. H and He have only the K shell, so they are period 1. Li through Ne have K and L, so they are period 2.
- The number of elements in a period depends on how the electrons fill. Each period begins with a new main shell's s sub-shell and ends when that shell is filled with respect to s and p — the first period excepted.
Working through it:
| Period | Sub-shells filled | Elements | Composition |
|---|---|---|---|
| 1 | 1s only | 2 | 1s¹ (H) and 1s² (He) |
| 2 | 2s, 2p | 8 | 2 s-block (Li, Be) + 6 p-block (B to Ne) |
| 3 | 3s, 3p | 8 | 2 s-block (Na, Mg) + 6 p-block (Al to Ar) |
| 4 | 4s, 3d, 4p | 18 | 2 s-block (K, Ca) + 10 d-block (Sc to Zn) + 6 p-block (Ga to Kr) |
| 5 | 5s, 4d, 5p | 18 | ₃₇Rb to ₅₄Xe |
| 6 | 6s, 4f, 5d, 6p | 32 | 2 s + 14 f + 10 d + 6 p, ₅₅Cs to ₈₆Rn |
| 7 | 7s, 5f, 6d, 7p | incomplete | 2 s + 14 f + 10 d + some p |
The third period contains eight and not eighteen elements because although the M shell has a 3d sub-shell, 3d gets electrons only after 4s is filled. In the fourth period the order of entry is 4s, then 3d, then 4p.
The 4f elements are the lanthanoids and the 5f elements the actinoids, from ₉₀Th to ₁₀₃Lr. Both are shown separately at the bottom of the table.
The book's Do you know? box is honest about the boundary: some scientists take lanthanoids as ₅₇La to ₇₀Yb, some ₅₈Ce to ₇₁Lu, some ₅₇La to ₇₁Lu, and there is an argument for including ₂₁Sc and ₃₉Y as well — all of which have substance, because all of them have similar outer shell configurations. The suffix -ide means heir, as in Cl to Cl⁻, and -oid means the same.
8. Metals, Non-metals and Metalloids
The rule of thumb the chapter gives is by outer-shell electron count:
- Three or fewer electrons in the outer shell — metal.
- Five or more — non-metal.
- There are exceptions to both.
All s-block elements are metals. The d-block, groups 3 to 12, are metals, known as transition metals, and their metallic character decreases gradually from left to right. Lanthanoids and actinoids actually belong to group 3 (IIIB) within the transition elements, which is why they are called inner transition elements.
Metalloids or semi-metals have properties intermediate between metals and non-metals — they look metallic but are brittle like non-metals, and are generally semiconductors. Examples: B, Si, As, Ge.
In the p-block, except group 18, you find metals, non-metals and metalloids together. The table has a visible staircase-like demarcation: elements to its left are metals, to its right non-metals, and those on it or very near it are the metalloids.
9. Periodic Properties and Why They Trend
The modern table is organised on electronic configuration, and physical and chemical properties follow from the outer shell configuration. Within a group the outer configurations are similar, so chemical properties are similar and physical properties grade regularly from top to bottom. Across a period the atomic number rises by one each step, so no two elements share a valence configuration — properties differ, with regular gradation left to right.
Valency
Valency was defined as the combining power of an element with respect to hydrogen or oxygen. With respect to hydrogen it is the number of hydrogen atoms one atom combines with; with respect to oxygen it is twice the number of oxygen atoms.
One Na atom combines with one H to give NaH, so the valence of Na is 1. One Ca combines with one O to give CaO, so the valency of Ca is 2.
In general the valence with respect to hydrogen is the traditional group number, and for group V and above it is 8 − group number — chlorine gives 8 − 7 = 1. Note that the group number gives the maximum valence in the old I-to-VIII notation, but it is 18 − group number in IUPAC numbering.
Across a period the valency starts at 1, rises to 4, then falls 4, 3, 2, 1, 0 — for main-group elements with respect to hydrogen. Activity 3 asks you to find the valencies of the first 20 elements and describe how valency varies across a period and down a group.
Modern usage often takes valence as simply the number of valence-shell electrons, and the oxidation number concept has largely replaced valence in the current literature.
Atomic radius
The atomic radius is the distance from the centre of the nucleus to the outermost shell. It cannot be measured on an isolated atom, because the electron's location cannot be pinned down. So it is measured indirectly:
- Metallic radius — measure the distance between the nuclei of adjacent atoms in the solid and take half. More than 75 per cent of elements are metals, so this suits most of them.
- Covalent radius — measure the bond length between the two atoms in a covalent molecule such as Cl₂ and take half.
Atomic radius is measured in picometres, 1 pm = 10⁻¹² m.
Down a group it increases. The atomic number rises, more shells are needed to hold the extra electrons, and the distance from nucleus to outer shell grows.
Across a period it decreases. Electrons enter the same main shell — or an inner one, for d and f block elements — so the nucleus-to-outer-shell distance does not change, but nuclear charge increases with atomic number. The stronger pull on the outer electrons shrinks the atom.
Table-6 and Table-7 give the numbers, in pm:
| Set | Values |
|---|---|
| Group 1 | Li (152), Na (186), K (231), Rb (244), Cs (262) |
| Group 17 | F (64), Cl (99), Br (114), I (133), At (140) |
| Period 2 | Li (152), Be (111), B (88), C (77), N (74), O (66), F (64) |
| Period 3 | Na (186), Mg (160), Al (143), Si (117), P (110), S (104), Cl (99) |
Ions are not the same size as their atoms
Sodium has 11 protons and 11 electrons, with the outer electron 3s¹. Na⁺ has 11 protons but 10 electrons, and its 3s level is empty, so its outer configuration is 2s² 2p⁶. With more protons than electrons, the nucleus pulls the remaining electrons harder and the ion shrinks. A cation is always smaller than its neutral atom.
Chlorine has 17 protons and 17 electrons, 1s² 2s² 2p⁶ 3s² 3p⁵. Cl⁻ has 17 protons and 18 electrons, 1s² 2s² 2p⁶ 3s² 3p⁶. Nuclear attraction per electron is less, so the ion swells. An anion is always larger than its neutral atom.
| Cations (pm) | Anions (pm) |
|---|---|
| Na 157, Na⁺ 98 | Cl 99, Cl⁻ 181 |
| K 203, K⁺ 133 | F 64, F⁻ 133 |
| Mg 136, Mg²⁺ 65 | O 73, O²⁻ 140 |
| Al 125, Al³⁺ 50 | N 75, N³⁻ 171 |
The book asks which member of each pair is larger: (a) Na, Al (b) Na, Mg²⁺ (c) S²⁻, Cl⁻ (d) Fe²⁺, Fe³⁺ (e) C⁴⁻, F⁻.
Ionization energy
The ionization energy is the energy required to remove an electron from the outermost shell of a neutral gaseous atom. Removing the first electron gives the first ionization energy; removing one from the uni-positive ion gives the second, and so on:
M(g) + IE₁ -> M⁺(g) + e⁻
M⁺(g) + IE₂ -> M²⁺(g) + e⁻
Five factors decide it:
| Factor | Effect | The book's example |
|---|---|---|
| Nuclear charge | More charge, more energy needed | Cl has higher IE than Na |
| Screening or shielding effect | More inner shells screen the nucleus, lowering IE | Cs has lower IE than Li |
| Penetration power of the orbital | 4s > 4p > 4d > 4f, so 4f is easiest to remove | B has lower IE than Be |
| Stable configuration | Half-filled and filled shells resist removal | O loses an electron more easily than N |
| Atomic radius | Larger radius, lower IE | F has higher IE than I |
Ionization energy decreases down a group and generally increases from left to right across a period. It is expressed in kJ mol⁻¹; when the term ionization potential is used, the unit is better written eV atom⁻¹.
The book prints a full table of first ionization energies in kJ mol⁻¹, whose left-to-right rise and top-to-bottom fall you can check yourself — for example Li 520.2, Be 899.5, B 800.6, C 1086.5, N 1402.3, O 1313.9, F 1681, Ne 2080.7 across period 2, and Li 520.2, Na 495.9, K 418.8, Rb 403.0 down group 1.
Notice that the rise across period 2 is not monotonic: B dips below Be, and O below N, which is exactly the penetration and stable-configuration effect.
A second table gives successive ionization energies, and the pattern in it answers the Think and discuss question of why IE₂ always exceeds IE₁ — for lithium the values are 520.2, 7300 and 11750.
Electron affinity, or electron gain enthalpy
The electron affinity is the energy liberated when an electron is added to a neutral gaseous atom. It is also called the electron gain enthalpy.
M(g) + e⁻ -> M⁻(g) + E₁
M⁻(g) + e⁻ -> M²⁻(g) + E₂
Practically no element liberates energy when a second electron is added to its uni-negative ion. That does not mean di-negative and tri-negative ions do not form — they do, but the energy for the second electron has to come from somewhere else, such as bond formation.
Table-8 gives values in kJ mol⁻¹:
| Group | Electron affinity values |
|---|---|
| VIIA (halogens) | F (−328); Cl (−349); Br (−325); I (−295); At (−270) |
| VIA (chalcogens) | O (−141); S (−200); Ge (−195); Te (−190); PO (−174) |
The values decrease down a group and increase across a period. Metals have very low values and alkaline earth metals have positive ones. The negative sign means energy is liberated; a positive sign means energy is absorbed.
Two Think and discuss questions sit here, and both are answerable from the table: why the calculated values for alkaline earth metals and noble gases are positive, and why fluorine — a second-period element — has a smaller electron gain enthalpy than chlorine below it.
Electronegativity
Ionization energy and electron gain enthalpy are properties of isolated atoms. Electronegativity was introduced to compare the pull of atoms inside a bond: it is the relative tendency of an atom to attract electrons towards itself when bonded to an atom of another element.
Mulliken proposed that it be taken as the average of ionization energy and electron affinity:
electronegativity = (ionization energy + electron affinity) / 2
Pauling instead assigned values from bond energies, fixing hydrogen at 2.20 and computing the rest relative to it. Table-9 gives his values:
| Set | Values |
|---|---|
| Halogens | F (4.0), Cl (3.0), Br (2.8), I (2.5) |
| Period 2 | Li (1.0), Be (1.47), B (2.0), C (2.5), N (3.0), O (3.5), F (4.0), Ne (−) |
Electronegativity decreases down a group and increases across a period. The most electronegative element is F, and the least electronegative stable element is Cs.
Metallic and non-metallic character
Metals show low electronegativity and tend to exist as positive ions in compounds — this is electropositive character. Non-metals are more electronegative, because their atomic radii are smaller.
Take period 3: Na, Mg, Al, Si, P, S, Cl. Na and Mg are metals, Al and Si are metalloids, P, S and Cl are non-metals. So metals sit on the left and non-metals on the right, and metallic character decreases across a period.
Take group 14: C, Si, Ge, Sn, Pb. Carbon is a non-metal, Si and Ge are metalloids, Sn and Pb are metals. So metallic character increases down a group.
The trends in one table
This summary table is reproduced from the book's own end-of-chapter list:
| Periodic property | Down a group | Across a period (left to right) |
|---|---|---|
| Valency | Same for all elements | Increases from 1 to 4 and decreases to 0 |
| Atomic radius | Increasing | Decreasing |
| Ionisation energy | Decreasing | Generally increasing |
| Electron affinity | Decreasing | Increasing |
| Electronegativity | Decreasing | Increasing |
| Electropositivity | Increasing | Decreasing |
| Metallic nature | Increasing | Decreasing |
| Non-metallic nature | Decreasing | Increasing |
Key words from the chapter
Triad, octave, periodic law, periodic table, period, group, lanthanides, actinides, element family, metalloids, periodicity, atomic radius, ionization energy, electron affinity, electronegativity, electropositivity.
10. Summary
Classification became necessary as the count of elements grew from about 13 in Boyle's time to more than 118 today, and the old metal/non-metal split failed on cases like aluminium.
Dobereiner (1829) grouped elements into triads in which the middle atomic weight is the average of the outer two. The law covered only some elements, failed at very low and very high masses, and weakened as atomic masses were measured more accurately — but it established that weights correlate with properties.
Newlands (1865) proposed the law of octaves, that every eighth element repeats the properties of the first, by analogy with the musical scale. He was the first to number the elements. His table jammed two elements into one slot, grouped dissimilar elements together, stopped at 56 elements and left no room for discoveries.
Mendeleeff ordered elements by atomic weight in 8 groups and 7 periods, stated that properties are periodic functions of atomic weight, left blank spaces for undiscovered elements and predicted the properties of eka-aluminium and eka-silicon — later found as gallium and germanium.
He also corrected the atomic weights of Be, In and Au, and accepted anomalous pairs such as Te before I. Its limitations were exactly those anomalous pairs, and the placing of dissimilar elements in the same group as sub-groups A and B.
Moseley (1913) showed by X-ray analysis that the atomic number is more fundamental than the atomic weight, and the periodic law was restated on atomic number, and then on electronic configuration.
In the long form, 18 groups and 7 periods, an element's group follows from its valence shell configuration and its period from the number of main shells. Elements are s, p, d or f block according to which sub-shell the differentiating electron enters.
Period lengths of 2, 8, 8, 18, 18 and 32 follow from the filling order, in which 3d fills only after 4s. The d-block are transition elements and the f-block lanthanoids and actinoids are the inner transition elements.
Across the table, atomic radius falls and nuclear pull rises left to right, so ionization energy, electron affinity, electronegativity and non-metallic character all increase, while metallic and electropositive character decrease. Down a group, added shells increase the radius and screening, so the first four fall and the last two rise. Cations are smaller and anions larger than their parent atoms, because the proton-to-electron ratio changes.
