Chapter 8: Periodic Table and PeriodicityShort Questions with Answers

Atomic number is the number of protons, which is fixed and different for every element; it also equals the number of electrons in a neutral atom, which decides the electronic configuration and so the chemical properties. When elements are arranged by atomic number, elements with similar outer configurations fall in the same group and properties repeat regularly.

Atomic mass did not work perfectly: argon (39.9) is heavier than potassium (39.1) but must come before it; the same happens with Co/Ni and Te/I.

Periodic means repeating after regular intervals. When elements are arranged by increasing atomic number, similar properties recur after 2, 8, 8, 18, 18, 32 elements. This trend is called periodicity in the properties, and that is why the table is called periodic.

Each period represents the filling of a shell, and the number of elements equals the number of electrons the subshells filled in that period can hold. Higher shells have more subshells: period 2 fills 2s, 2p (8); period 4 adds 3d (2 + 10 + 6 = 18); period 6 also adds 4f (2 + 14 + 10 + 6 = 32). So periods get longer: 2, 8, 8, 18, 18, 32, 32.

Each new period starts filling a new shell. Going down a group, one more complete shell is added inside, but the outermost shell is filled in the same way each time: Li 2, 1; Na 2, 8, 1; K 2, 8, 8, 1. The pattern of outer electrons repeats, so the elements that line up in a column have the same number of outer electrons.

No. Sodium is more reactive. Sodium (group 1) has to lose only one electron (ionisation energy 496 kJ mol⁻¹), while calcium (group 2) has to lose two (590 + 1145 kJ mol⁻¹). Metallic character decreases from left to right in a period. Sodium reacts violently with cold water; calcium reacts steadily.

Radius decreases from left to right in a period. Sodium (186 pm) has the maximum radius. Among the elements that form bonds, chlorine (99 pm) has the minimum.

The "sixth and seventh groups" are groups 16 and 17 (old VIA and VIIA). Their atoms have 6 or 7 outer electrons, a large nuclear charge and a small size for their period, so they attract a shared pair strongly. They need only 1 or 2 electrons to complete the octet. Fluorine (4.0), oxygen (3.4) and chlorine (3.2) are the most electronegative elements.

The first electron is removed from a neutral atom; the second must be removed from a positive ion Mg⁺, in which the same 12 protons hold one electron fewer, so the remaining electrons are held more tightly. More energy is needed.

Generally no. An ionic bond needs one atom that loses electrons easily (a metal, low ionisation energy) and one that gains them (a non-metal, high electron affinity), with a large difference in electronegativity. Two metals both want to lose electrons: they form a metallic bond (alloys). Two non-metals both want to gain: they share electrons (covalent).

Least ionisation energy: caesium (377 kJ mol⁻¹; francium would be lower still, but it is radioactive and rare), at the bottom of group 1. Highest electronegativity: fluorine (4.0), at the top of group 17.

s²p³ = 5 outer electrons → group 15 (nitrogen family), last subshell p → p-block. The period equals the number n of the outermost shell, which "s²p³" does not give: 2s²2p³ is nitrogen (period 2), 3s²3p³ phosphorus (period 3).

Group 18, the noble gases (ns²np⁶; except helium which has 2). They exist as gases made of single atoms: Ne, Ar, Kr, Xe, Rn.

The only gas in group 16 is oxygen, in period 2 (2, 6). Sulphur, selenium and tellurium are solids.

Radius increases down a group and decreases across a period, so the largest atom is at the bottom left: group 1, period 7 (francium). Of the common stable elements, caesium (group 1, period 6) is largest.

Most electropositive: group 1 (Cs, Fr). Most electronegative: group 17 (fluorine).

Lowest ionisation energy: K (419 kJ mol⁻¹), group 1, period 4 (largest atom of the five, on the left). Highest electron affinity: Cl (−349 kJ mol⁻¹), group 17, period 3.

Metals give basic oxides; non-metals give acidic oxides.

The horizontal rows of elements are called periods while the vertical columns are called groups. There are 7 periods and 18 groups.

The repetition of the properties of elements after regular intervals is called periodicity. It occurs because similar outer electronic configurations recur.

Mendeleev arranged elements in ascending order of atomic masses; Moseley arranged them in ascending order of atomic numbers. The modern table follows Moseley.

A period begins when a new shell receives its first electron (ns¹: an alkali metal) and ends when that shell's s and p subshells are full (ns²np⁶: a noble gas). Period 1 starts with hydrogen (1s¹), which is not an alkali metal, and ends with helium.

Two series of fourteen elements in the 6th and 7th periods: lanthanides after lanthanum (57) and actinides after actinium (89). They are placed at the bottom because of the space problem, to keep the table manageable. Together they are the f-block or rare earths.

Normal elements (groups 1, 2 and 13-17): all inner shells complete, outermost incomplete. Transition elements (groups 3-12): inner sub-shells are in the process of completion.

s, p, d and f blocks, depending on the type of subshell being filled. s: groups 1-2; p: groups 13-18; d: groups 3-12; f: the two bottom rows.

Group 1: alkali metals; group 2: alkaline earth metals; group 17: halogens; group 18: noble gases. Groups 13, 14, 15, 16 are the boron, carbon, nitrogen and oxygen families.

2, 8, 6 (1s² 2s² 2p⁶ 3s² 3p⁴): three shells → period 3; last subshell p → p-block; 6 outer electrons → group 16.

The chemical properties depend largely upon the number of electrons in the outermost shell. All elements of a group have the same number of electrons in the outermost shell, so they show similar chemical properties.

They have one electron in their outermost shells, so they show a strong tendency to lose it, forming cations. They react vigorously with water giving hydrogen and alkalis.

Down the group the atoms get bigger and the outer electron is further from the nucleus and better shielded, so it is lost more easily (lower ionisation energy). Li reacts gently with water, Na vigorously, K catches fire.

They are very reactive non-metals that form salts with metals; hence they are called halogens (salt forming). Na + ½Cl₂ → NaCl.

Halogens react by gaining an electron. Down the group the atoms get bigger, so the nucleus attracts an incoming electron less strongly: F > Cl > Br > I. Unlike alkali metals, the reactivity of halogens decreases down the group.

Main-group metals form cations with a charge equal to the group number (group 1 → 1+, group 2 → 2+, group 13 → 3+), equal to their outer electrons. Non-metals form anions with a negative charge equal to the number of electrons needed to complete the octet: group 17 → 1−, group 16 → 2−, group 15 → 3−.

They have s²p⁴ configuration, so they have a tendency to accept two electrons to form a dinegative ion: O + 2e⁻ → O²⁻.

Atomic radius is half the distance between the nuclei of two identical bonded atoms. It is expressed in picometres (1 pm = 10⁻¹² m). Carbon: 154 ÷ 2 = 77 pm (covalent radius).

The electron cloud of an atom has no definite limit. So the radius is measured when the atom is bonded to an identical atom.

Electrons are added to the same shell while the nuclear charge increases (Li +3 to Ne +10). The greater pull draws the electrons closer, so atoms shrink: Li 152 pm to F 71 pm.

A new shell is added in each successive period, and the inner electrons shield the outermost electrons from the nucleus. Li 152, Na 186, K 227, Rb 248, Cs 265 pm.

Na > Mg (same period; Mg has a larger nuclear charge). K > Na (K is lower in group 1, one more shell).

Ionisation energy is the amount of energy required to remove the most loosely bound electron from the valence shell of an isolated gaseous atom. Na(g) → Na⁺(g) + e⁻, +496 kJ mol⁻¹.

The smaller the radius of an atom, the stronger the attraction between the nucleus and the outer electrons, and the higher the ionisation energy.

Nuclear charge increases and atomic size decreases, so the outer electrons are held more tightly: Li 520 to Ne 2081 kJ mol⁻¹.

The outer electron is in a higher shell, further from the nucleus and shielded by more inner electrons, so it is removed more easily: Li 520, Na 496, K 419, Rb 403, Cs 377 kJ mol⁻¹.

They have complete outer shells (ns²np⁶) and are the smallest atoms of their periods, with the largest nuclear charge, so their electrons are held very tightly: He 2372, Ne 2081 kJ mol⁻¹. This is why they hardly react.

Electron affinity is the amount of energy change when an electron is added to the outermost shell of an isolated gaseous atom. F(g) + e⁻ → F⁻(g), ΔH = −328 kJ mol⁻¹.

The sizes of atoms increase down the group. In a bigger atom the nucleus attracts the incoming electron with a weaker force, so the electron affinity is lower.

Fluorine is very small, so its seven outer electrons are crowded close together; an extra electron is repelled by them, which partly cancels the strong attraction of the nucleus. Chlorine is bigger, with more room, so it releases more energy (−349 against −328 kJ mol⁻¹).

The second electron is added to a negative ion (O⁻ + e⁻ → O²⁻), which repels it, so energy must be supplied: the value is positive.

The force with which an atom attracts the shared pair of electrons towards itself in a bond. It increases from left to right in a period and decreases from top to bottom in a group. Li 1.0 ... F 4.0; F 4.0, Cl 3.2, Br 3.0, I 2.7.

The most electronegative atoms are found at the top right-hand corner (F, O, N, Cl). The most electropositive elements lie at the bottom left of the periodic table (Cs, Fr).

The tendency of an element to lose electrons and form cations. It increases down a group (ionisation energy falls) and decreases across a period (nuclear charge rises).

It decreases from the very reactive sodium to magnesium, aluminium and silicon, as more electrons must be lost; then it increases again towards the non-metals phosphorus, sulphur and chlorine, which react by gaining electrons. Argon is unreactive.

Density generally increases from top to bottom in a group: Li 0.53, Na 0.97, Rb 1.53, Cs 1.93 g cm⁻³. Potassium (0.86) is an exception: lighter than sodium. Li, Na and K float on water.

Each alkali metal atom gives only one electron to the electron sea, and the atoms are large, so the metallic bonding is weak. Down the group the atoms get even bigger and the bonding weaker, so the melting points fall: Na 98, K 63, Rb 39, Cs 28 °C.