Periodic Table & Periodicity

With more than a hundred elements known, studying each one separately is impossible, so chemists arranged them into a table that groups elements of similar nature together.

• Elements with similar properties are placed together, so learning one helps predict the rest.
• The arrangement is called the periodic table, because properties repeat at regular (periodic) intervals.
• The repetition of properties at fixed intervals is called periodicity.

The earliest serious attempt grouped elements in sets of three, after the German chemist Johann Dobereiner noticed a neat numerical pattern among them.

• The atomic mass of the middle element was nearly the average of the other two.
• Example: in the triad Li, Na, K, the atomic mass of Na (23) ≈ average of Li (7) and K (39) = (7 + 39)/2 = 23.
• Other triads: Ca, Sr, Ba and Cl, Br, I.
• Limitation: only a few elements could be arranged in triads, so the idea failed for most elements.

The English chemist John Newlands lined up elements by increasing atomic mass and found that every eighth element repeated the properties of the first, like notes in music.

• Every eighth element had properties similar to the first, like the 8 notes of a musical octave (sa-re-ga-ma).
• Example: starting from Li, the eighth element is Na, which resembles Li.
• Limitation: it worked only up to calcium (Ca); beyond it the pattern broke down.
• Limitation: no gaps were left for undiscovered elements, and dissimilar elements were forced together.

The Russian chemist Dmitri Mendeleev produced the first widely accepted periodic table, and his bold predictions about missing elements made it a triumph.

• Elements were arranged in increasing order of atomic mass.
• The table had vertical columns called groups and horizontal rows called periods.

• Mendeleev left gaps for undiscovered elements rather than forcing the order.
• He predicted the properties of unknown elements, naming them eka-aluminium, eka-boron and eka-silicon.
• These were later discovered as gallium, scandium and germanium, matching his predictions closely.
• He even corrected some atomic masses (e.g. beryllium) to fit the pattern.

Despite its success, Mendeleev's table had clear flaws, most of which came from using atomic mass as the basis of arrangement.

• Isotopes had no place, since atoms of the same element with different masses would need different positions.
• Some heavier elements came before lighter ones (anomalous pairs), e.g. argon (Ar, 40) was placed before potassium (K, 39).
• No place was given for the noble gases (they were not yet discovered).

The English physicist Henry Moseley showed that the real basis of order is not mass but the number of protons, and this fixed almost every defect of the older table.

• Atomic number (Z) = number of protons in the nucleus, and was found to be the fundamental property.
• Arranging by atomic number removed the anomalous pairs (Ar before K now makes sense, as Z of Ar = 18 < Z of K = 19).
• Isotopes fit in one place, since all isotopes of an element have the same atomic number.

The two laws look alike but differ in one decisive way — the property used to arrange the elements.

The modern table is a long-form grid built from vertical groups and horizontal periods, into which all the elements fall in a fixed order of atomic number.

• There are 7 horizontal rows called periods.
• Elements in the same group have similar chemical properties, because they have the same number of outermost (valence) electrons.
• The period number = number of electron shells in the atom.

Depending on which sub-shell the last electron enters, the table is divided into four blocks — s, p, d and f.

• p-block — groups 13 to 18; contains metals, non-metals and metalloids.
• d-block — groups 3 to 12; the transition metals.
• f-block — the lanthanides and actinides, placed in two rows at the bottom (also called inner transition elements).

Certain groups behave so distinctively that they are given special family names, and these are favourites in exams.

The first two groups are the most reactive metals, so reactive that they are never found free in nature.

• They are the most reactive metals and are stored under kerosene oil.
• They are soft and can be cut with a knife (e.g. sodium, potassium).
• Hydrogen is placed in group 1 but is a non-metal, not an alkali metal.

• They have 2 electrons in their outermost shell.
• They are less reactive than alkali metals but still quite reactive.
• Calcium and magnesium are common examples, important in bones and chlorophyll.

The right edge of the table holds the most reactive non-metals and, beyond them, the famously unreactive noble gases.

• They are the most reactive non-metals.
• Fluorine is the most reactive of all non-metals.
• The word "halogen" means salt-producer, as they form salts (e.g. NaCl) with metals.

• They have a complete (stable) outermost shell — 2 electrons for He, 8 for the rest.
• They are chemically inert (unreactive), hence called inert or noble gases.
• Helium is the lightest noble gas; argon is the most abundant noble gas in air.

The table also splits, left to right, into metals, non-metals and a thin diagonal band of in-between elements.

• Non-metals are on the upper-right side of the table.
• Metalloids lie along the zig-zag "staircase" separating metals from non-metals.

• Metalloids show properties of both metals and non-metals.
• Examples: boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te).
• Silicon and germanium are used as semiconductors in electronics.

Atomic size changes in an orderly way across the table, shrinking along a period and growing down a group.

• Down a group (top to bottom), atomic size increases, as new electron shells are added.
• The largest atom is at the bottom-left; the smallest is at the top-right of the table.

Whether an element behaves as a metal or a non-metal also follows a clear direction across the table.

• Down a group, metallic character increases (and non-metallic character decreases).
• The most metallic elements are at the bottom-left (e.g. caesium, francium).
• The most non-metallic elements are at the top-right (e.g. fluorine, excluding noble gases).

Reactivity depends on how easily an element loses or gains electrons, and it moves in opposite directions for metals and non-metals.

• Metallic reactivity decreases across a period from left to right.

• Down group 17, the reactivity of halogens decreases (F is more reactive than I).
• Non-metallic reactivity increases across a period from left to right.
• Fluorine is the most reactive non-metal; caesium/francium the most reactive metal.