Chapter 8Periodic Table and Periodicity

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.

First find its atomic number (from its X-ray spectrum, as Moseley did) and its electronic configuration. Then:

A new element heavier than oganesson (118) would have Z = 119 and configuration 8s¹: it would start period 8 in group 1, below francium.

Hydrogen, Z = 1, configuration 1s¹. It can do both:

That is why hydrogen is placed at the top of group 1 but is not an alkali metal.

Aluminium. Both are in group 13 with three outer electrons, but aluminium is bigger (one more shell). Its outer electrons are further from the nucleus and more shielded, so they are held less strongly: the smaller the atom, the higher the ionisation energy.

So aluminium (a metal, forms Al³⁺) loses electrons more easily than boron (a metalloid that forms covalent compounds). Metallic character increases down the group.

The electron cloud of an atom has no definite limit, so its size cannot be defined exactly. It is possible to measure the radius of an atom when it is bonded to an identical atom.

The distance between the nuclei of two bonded carbon atoms is 154 pm. Half of this, 77 pm, is the radius of the carbon atom, also called its covalent radius. For metals, half the distance between neighbouring atoms in the crystal is used (metallic radius).

Oxygen (2, 6) has six outer electrons and needs only two to complete its octet: O + 2e⁻ → O²⁻ (2, 8), the configuration of neon. Nitrogen (2, 5) needs three: N³⁻.

A third electron for oxygen would have to go into a new, third shell, far from the nucleus, and would be strongly repelled by the already negative O²⁻ ion. No energy is gained by this, so O³⁻ does not form. Atoms gain only as many electrons as complete their octet.

Information needed: electronic configuration

The position of an element depends on its electronic configuration, which comes from its atomic number. If we know the configuration (or the number of shells and of outer electrons), we can place it:

Chemical properties help too

The charge on its ion, the nature of its oxide (basic or acidic), whether it is a metal or non-metal and its reactivity with water all point to its group.

Is atomic mass helpful?

Not reliably. Mendeleev used atomic mass, but in some pairs the heavier element must come first because of its properties:

Isotopes also have different masses but the same position. So the atomic number, not the mass, decides the position; atomic mass gives only a rough idea.

Four blocks

Elements in the periodic table are classified into blocks depending upon the type of the subshell which is being filled: s, p, d and f.

Yes, they are helpful

So knowing the block tells us at once the general behaviour of an element.

(a) Atomic radius

Atomic radius is half the distance between the nuclei of two identical bonded atoms. In a period it decreases from left to right.

Reason: as we go from lithium to neon, electrons are added to the same outermost shell, while the charge on the nucleus increases from +3 to +10. This pulls the electrons closer to the nucleus, so the atoms get smaller.

(b) Ionisation energy

Ionisation energy is the energy required to remove the most loosely bound electron from an isolated gaseous atom. In a period it increases from left to right.

Reason: the atoms get smaller and the nuclear charge larger, so the attraction between the nucleus and the outer electrons is stronger and more energy is needed to remove one.

Exceptions

B is lower than Be (a new 2p subshell starts) and O lower than N (O has its first paired 2p electron), but the overall trend is upward.

1. Ionisation energy

Elements with low ionisation energy (metals of groups 1 and 2) lose electrons easily and form positive ions: they form ionic bonds with non-metals, and metallic bonds with each other.

2. Electron affinity

Elements with high electron affinity (halogens, oxygen) gain electrons easily and form negative ions with metals.

3. Electronegativity

The difference in electronegativity between two atoms decides the bond:

4. Atomic size

Large atoms on the left lose electrons easily (ionic, metallic); small atoms on the right attract electrons (covalent or anions).

5. Metallic character

Metal + metal → metallic bond; metal + non-metal → ionic; non-metal + non-metal → covalent.

Boron and silicon are also solids, but they are metalloids. Bromine is the only non-metal that is a liquid at room temperature; the rest (H₂, N₂, O₂, F₂, Cl₂, noble gases) are gases.

A period fills the outermost s and p

Each period represents the completion of a shell, and the number of elements in a period is fixed by the number of electrons accommodated. The elements of a period run from ns¹ to ns²np⁶.

Periods 2 and 3

Period 2 fills 2s and 2p: 2 + 6 = 8. Period 3 fills 3s and 3p: 2 + 6 = 8. The third shell also has a 3d subshell, but 3d is higher in energy than 4s, so it is filled only in period 4 (from scandium). So periods 2 and 3 both have 8 elements.

Other periods

Period 1 has only the 1s subshell (2 elements); periods 4 and 5 add ten d-block elements (18); periods 6 and 7 also add fourteen f-block elements (32).

Benefits

So the periodic table is the most useful single summary of chemistry.

1. Very small size

Li and Be are the first members of their groups, with only two shells; they are the smallest atoms of their groups (Li 152 pm, Be 113 pm).

2. High charge density of the ions

Li⁺ and especially Be²⁺ are tiny with a concentrated charge, so they pull the electron clouds of neighbouring anions and give their compounds covalent character (BeCl₂ is largely covalent).

3. High ionisation energy and electronegativity

They hold their electrons more strongly (Li 520, Be 899 kJ mol⁻¹), so they are less electropositive and less reactive than the members below them.

4. No d-orbitals in the outer shell

Their second shell has only s and p subshells, so Be can hold at most four bonds.

Examples of different behaviour

Mendeleev's merits

He grouped elements with similar properties, left gaps and predicted the properties of undiscovered elements (gallium, scandium, germanium), and corrected some atomic masses.

Why the modern table is better

Using atomic number removes the anomalies, explains isotopes, and links the position of every element directly to its electronic configuration.

Basis

The modern periodic table is based upon the arrangement of elements according to increasing atomic number. Properties of elements repeat after regular intervals, so elements of similar properties are placed in the same vertical columns.

Periods

There are seven periods. Each period except the first starts with an alkali metal and ends at a noble gas. Each period represents the completion of a shell; the number of elements in a period is fixed.

Lanthanides and actinides

In the 6th and 7th periods two series of fourteen elements each are placed at the bottom: lanthanides (after La = 57) and actinides (after Ac = 89). Together they are called rare earths or f-block elements.

Groups

There are eighteen groups. Elements in a group resemble one another in chemical properties since they contain the same number of electrons in their outermost shells. A group is also called a family.

Normal and transition elements

Groups 1, 2 and 13 to 17 contain normal elements: inner shells complete, outermost shell incomplete. Groups 3 to 12 are transition elements: inner sub-shells are being completed.

Blocks

s-block (groups 1, 2), p-block (13-18), d-block (between s and p) and f-block (two rows at the bottom), according to the subshell being filled.

Reason

The chemical properties of elements depend largely upon the number of electrons in their outermost shells; all elements of a group have the same number, so they show similar properties.

Group 1 (ns¹)

Strong tendency to lose their one electron, forming cations: electropositive metals. They react vigorously with water producing hydrogen and alkalis. Their oxides are strongly basic; they form halides with halogens. Reactivity increases down the group.

Group 2 (ns²)

They lose both outer electrons, forming dipositive ions; the tendency increases down the group. Their oxides are basic and form alkalis in water.

Group 16 (ns²np⁴)

They accept two electrons to form dinegative ions; oxygen is strongly electronegative. O + 2e⁻ → O²⁻; they form oxides and sulphides.

Group 17 (ns²np⁵)

A strong tendency to gain one electron to become an anion; very reactive non-metals that form salts with metals (halogens). Reactivity decreases down the group.

Charge on ions

Ionisation energy

The energy required to remove the most loosely bound electron from the valence shell of an isolated gaseous atom. Na → Na⁺ + e⁻, +496 kJ mol⁻¹. It increases from left to right in a period and decreases from top to bottom in a group, because smaller atoms hold their electrons more strongly.

Electron affinity

The energy change when an electron is added to the outermost shell of an isolated gaseous atom. F + e⁻ → F⁻, −328 kJ mol⁻¹. Smaller atoms attract the entering electron more strongly, so it increases across a period and decreases down a group (Cl −349 is an exception, above F).

Electronegativity

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 down a group. Most electronegative atoms are at the top right (F, O, N, Cl).

Difference between EA and EN

Electron affinity is an energy change of a free gaseous atom gaining an electron; electronegativity is the pull of a bonded atom on a shared pair and has no unit.

Metallic character

The tendency of an element to lose electrons and form positive ions (cations).

In a group

Since ionisation energy decreases down a group, elements have an increased ability to lose electrons, so both metallic character and reactivity increase down the group. Group 2: Ba > Sr > Ca > Mg. Group 14 goes from non-metal carbon to metals tin and lead.

In a period

Moving left to right the nuclear charge increases and the valence electrons are pulled strongly, making it difficult to lose them, so metallic character decreases. Period 3: Na, Mg, Al are metals; Si a metalloid; P, S, Cl non-metals.

Reactivity in a period

Reactivity decreases from the very reactive sodium to magnesium, aluminium and silicon; moving further right towards the non-metals, reactivity increases again (chlorine is very reactive).

Density

Density is the mass of a substance in unit volume. It generally increases from top to bottom in a group but varies less across a period.

Rules

Example 1: sodium, Z = 11

1s² 2s² 2p⁶ 3s¹ → period 3, s-block, group 1.

Example 2: chlorine, Z = 17

1s² 2s² 2p⁶ 3s² 3p⁵ → period 3, p-block, group 10 + 7 = 17.

Example 3: calcium, Z = 20

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² → period 4, s-block, group 2.

Example 4: iron, Z = 26

… 3s² 3p⁶ 3d⁶ 4s² → period 4, d-block, group 6 + 2 = 8.

Example 5: neon, Z = 10

1s² 2s² 2p⁶ → period 2, p-block, group 18.