Chapter 1: Periodic Table, Periodic Properties and Variations of Properties
In 1913, a young English physicist named Henry Moseley set up an unusual experiment. He fired fast-moving electrons at plates of different metals — calcium, titanium, chromium, iron, cobalt, nickel, copper, zinc — and measured the X-rays that each metal gave out.
Every metal gave a slightly different X-ray. When Moseley plotted his measurements, the points did not scatter. They fell along a straight line, and the line climbed in neat whole-number steps as he moved from one metal to the next.
Those whole numbers were the count of protons in the nucleus. Moseley had found a way to weigh nothing at all and still put the elements in an exact order.
That order settled an old argument. Chemists had long noticed that cobalt behaves like a member of the iron family, even though cobalt is slightly heavier than nickel, which follows it. Arranged by mass, the two elements sat in the wrong places. Arranged by proton count, they sat correctly, and so did tellurium and iodine.
Moseley’s line also had visible gaps. Numbers 43, 61, 72 and 75 belonged to elements nobody had yet found. Each one was found later, exactly where the gap said it should be.
Moseley never saw that. He joined the army and was killed in 1915, aged twenty-seven. He was two years past the experiment that gave chemistry its modern table.
Diagram
Fig. 1.0-1 — Moseley’s Measurement Of Atomic Number Two panels side by side. Left panel: simple X-ray tube — heated filament emitting electrons, accelerating voltage marked, electron beam striking a slanted metal target, X-rays leaving the target towards a crystal and detector; label filament, electron beam, metal target, X-rays, detector. Right panel: a graph with atomic number on the horizontal axis (marked 20 to 30) and the square root of X-ray frequency on the vertical axis; plot a straight rising line through evenly spaced points, each point labelled with an element symbol in order Ca, Ti, Cr, Fe, Co, Ni, Cu, Zn; caption arrow pointing to the even spacing labelled ‘equal steps of one proton’.
1.1 Classifying The Elements
An element is a pure substance made up of only one kind of atom. Elements are the building blocks of every material around you — the copper in a wire, the oxygen in the air, the calcium in a bone.
More than a hundred elements are known today, and each one has its own set of properties. Studying them one at a time would be endless work. So chemists looked for a way of arranging them in an organised manner, placing elements that behave alike close together.
You met three of these early attempts in Class IX. It is worth walking through them again, because each one got a little closer to the arrangement we use now.
1.1.1 Early Attempts At Classification
The earliest arrangements used the only number chemists could measure at the time — the atomic weight of an element, which tells us how heavy one of its atoms is compared with a standard atom.
Döbereiner’s Triads
In the early nineteenth century, Johann Döbereiner noticed that certain elements fall naturally into sets of three with similar properties. He called such a set a triad — a group of three elements with similar properties, in which the atomic weight of the middle element is the arithmetic mean of the other two.
Two of his triads were:
\[^{40}_{20}\mathrm{Ca},\quad ^{88}_{38}\mathrm{Sr},\quad ^{137}_{56}\mathrm{Ba}\]
\[^{35}_{17}\mathrm{Cl},\quad ^{80}_{35}\mathrm{Br},\quad ^{127}_{53}\mathrm{I}\]
Take the calcium–strontium–barium set. The mean of the two outer atomic weights is \((40 + 137)\div 2 = 88.5\), which is very close to the atomic weight of strontium. The trouble was that only a handful of elements could be fitted into triads at all.
Newlands’ Octaves
Later, John Newlands arranged the known elements in increasing order of their atomic weights. He observed that every eighth element, beginning from any element, resembles the first in its physical and chemical properties.
This repetition after eight reminded him of the eight notes of a musical scale, so the pattern became known as the law of octaves. It worked for the lighter elements and then broke down completely for the heavier ones.
Mendeleev’s Periodic Table
Dmitri Mendeleev gave the first true periodic table of elements, based on a law of his own.
This arrangement let Mendeleev place the 63 elements known at that time in vertical columns, which he called groups, and in horizontal rows, which he called periods. Elements in the same column shared similar properties.
Mendeleev’s table did something no list had done before: it made predictions. Where no known element fitted a place, he left the place empty and described the element that would one day fill it.
Even so, the table had defects. It could not explain the positions of certain elements, the rare earth metals or the isotopes.
The isotopes were the deepest problem of the three. Isotopes are atoms of the same element with different masses, so ordering by mass demanded that they be given separate places — yet chemically they are one and the same element. A property that splits an element in two cannot be the property the table is built on.
1.1.2 Atomic Number And The Modern Periodic Law
The defects of Mendeleev’s table were removed when Henry Moseley put forward the modern periodic table, which was based on atomic number rather than atomic weight.
The atomic number of an element is the number of protons in the nucleus of its atom. In a neutral atom the positive protons are exactly balanced by negative electrons, so the atomic number also equals the number of electrons in the shells.
Now think about what actually happens in a chemical reaction. Atoms lose, gain or share electrons; the nucleus is never touched. The physical and chemical properties of an element therefore depend on the number of electrons and on their arrangement in the shells — and the number of electrons is fixed by the atomic number.
Atomic number is also unique. No two elements have the same atomic number, whereas two different atoms can easily have the same mass. Hence atomic number is the fundamental property that identifies an element.
Moseley restated the periodic law in terms of this new quantity.
The word ‘periodic’ is doing real work in that sentence. It signifies that if the elements are arranged in the order of their increasing atomic numbers, elements with similar properties will be repeated after regular intervals.
1.1.3 The Long Form Of The Modern Periodic Table
Later on Niels Bohr gave the extended form of the table, known as the long form of the modern periodic table. This is the table printed in laboratories and in this book, and it is shown in Fig. 1.1-1.
Diagram
Fig. 1.1-1 — The Long Form Of The Modern Periodic Table Full-page landscape chart of the modern periodic table, 18 groups across and 7 periods down, with the lanthanide and actinide series set out as two separate rows below the main body. Each box carries the atomic number at the top left, the symbol in large type, and the name below it. Number the groups both ways: new IUPAC numbers 1 to 18 along the top, old notation (IA, IIA, IIIB … VIIB, VIII, IB, IIB, IIIA … VIIA, 0) immediately beneath. Colour-band and key the following as a legend: active metals (groups 1 and 2), transition metals, heavy metals, metalloids, non-metals, halogens, noble gases. Mark the s-block, p-block, d-block and f-block along the base. Show a small inset box explaining the layout of one element’s box, using carbon as the model, with the state (solid, liquid, gas) indicated by symbol.
You will meet two sets of group numbers on such a chart, so it is worth knowing where they come from. The older notation used Roman numerals with the letters A and B, and a zero group for the noble gases. The International Union of Pure and Applied Chemistry (IUPAC) has recommended that the groups simply be numbered from 1 to 18 instead. Table 1.1 sets the two side by side.
Table 1.1 — Old and new notation for the groups of the periodic table
| Old notation | IA | IIA | IIIB | IVB | VB | VIB | VIIB | VIII | VIII | VIII | IB | IIB | IIIA | IVA | VA | VIA | VIIA | 0 |
|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
| New notation | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 |
Notice that the old scheme gave one name, VIII, to three columns at once. The new scheme gives every column its own number, which is one reason it is now preferred.
1.2 Salient Features Of The Modern Periodic Table
Look along the top of Fig. 1.1-1 and then down its left edge. The whole table is built from these two directions — eighteen columns crossing seven rows. Almost everything you will learn in this chapter follows from what those two directions mean.
1.2.1 Groups
The modern periodic table has eighteen vertical columns. These vertical columns are known as groups, and they are arranged from left to right in the order of increasing atomic number.
Each vertical column accommodates elements with the same number of electrons in the outermost shell. The outermost shell of an atom is called its valence shell, and the electrons in it are the valence electrons — the electrons an atom actually uses when it reacts.
That single rule is why a column is worth naming at all. Elements sharing a column share their valence electrons, and so they share their chemistry.
1.2.1.1 The Individual Groups And Their Families
Several groups are used so often that they carry names of their own. Fig. 1.2-1 shows where each of them sits.
Diagram
Fig. 1.2-1 — Families Of The Modern Periodic Table Simplified outline grid of the periodic table (18 columns, 7 rows) with no element symbols, only shaded and labelled regions. Shade and label: group 1 ‘alkali metals (except hydrogen)’; group 2 ‘alkaline earth metals’; groups 3 to 12 as one broad band ‘transition elements’; group 13 ‘boron family’; group 14 ‘carbon family’; group 15 ‘nitrogen family’; group 16 ‘oxygen family or chalcogens’; group 17 ‘halogens’; group 18 ‘noble gases’. Show hydrogen’s box at the top of group 1 in a contrasting shade with a small tag ‘a non-metal’. Below the main grid, show two detached rows labelled ‘lanthanides — group 3, period 6’ and ‘actinides — group 3, period 7’, joined to the group 3 box of periods 6 and 7 by thin lines. Label groups 1, 2 and 13 to 18 collectively along the base as ‘main group (representative) elements’.
- Group 1 — these elements are known as alkali metals (except hydrogen), as they form strong alkalis with water.
- Group 2 — the alkaline earth metals. They form weaker alkalis as compared to the group 1 elements.
- Groups 3, 4, 5, 6, 7, 8, 9, 10, 11 and 12 — these are known as the transition elements. All the elements of these groups are metals, and they have their two outermost shells incomplete.
- Group 13 — the boron family. Boron is the first member of the group.
- Group 14 — the carbon family.
- Group 15 — the nitrogen family.
- Group 16 — the oxygen family, also known as the chalcogens, meaning ore forming.
- Group 17 — these elements form salts and so are known as halogens, meaning salt formers.
- Group 18 — the zero group. Its elements are called noble gases or inert gases.
The noble gases deserve a second look, because their behaviour explains a great deal of what follows. These elements have their outermost orbit complete. Due to this stable electronic configuration, they hardly react with other elements.
1.2.1.2 Main Group, Transition And Inner Transition Elements
The elements of groups 1 and 2, and of groups 13 to 18, are known together as the main group elements, or representative elements, or normal elements. They are the elements whose behaviour the trends of this chapter describe most simply.
Their valence electrons follow a straightforward count. The elements of groups 1, 2, 13, 14, 15, 16 and 17 have 1, 2, 3, 4, 5, 6 and 7 electrons in their valence shell respectively. The outermost shell of all the elements of the main group is incomplete, except group 18.
Two rows sit apart, below the main body of the table. These are the lanthanides, or rare earths, which belong to group 3 of the sixth period, and the actinides, the radioactive elements, which belong to group 3 of the seventh period. Both sets are shown separately for a purely practical reason: they are large in number, and placing them inside the main body would distort the shape of the table.
1.2.2 Periods
There are seven horizontal rows in the modern periodic table. They are known as periods.
Each period begins with an element having one electron in its valence shell, and ends with an element whose outermost shell is completely filled. The number of shells present in an atom determines its period. Elements of period one have one shell, elements of period two have two shells, elements of period three have three shells, and so on.
The first period is a special case of the rule. It holds only hydrogen and helium, and helium’s single shell is full with two electrons, so the period still ends with a completely filled outermost shell.
The shell diagrams in Fig. 1.2-2 show the same two atoms drawn out, so you can see the shells being counted.
Diagram
Fig. 1.2-2 — Reading Period And Group From Shell Diagrams Two shell (Bohr) diagrams side by side, drawn to the same scale. Left: sodium atom, nucleus marked ‘11 p’, three concentric circles carrying 2, 8 and 1 electrons shown as dots; brace across the three circles labelled ‘3 shells → period 3’; arrow to the single outer dot labelled ‘1 valence electron → group 1’. Right: calcium atom, nucleus marked ‘20 p’, four concentric circles carrying 2, 8, 8 and 2 electrons; brace labelled ‘4 shells → period 4’; arrow to the two outer dots labelled ‘2 valence electrons → group 2’.
1.2.2.1 Length Of The Periods
The periods are not all the same length. The first is the shortest, holding two elements; the sixth and seventh are the longest, holding thirty-two each. Table 1.2 sets out the seven periods with the atomic numbers they cover.
Table 1.2 — The seven periods of the modern periodic table
| Period number | Type of period | Number of elements | Atomic numbers |
|---|---|---|---|
| 1 | Shortest | 2 | 1 and 2 |
| 2 | Short | 8 | 3 to 10 |
| 3 | Short | 8 | 11 to 18 |
| 4 | Long | 18 | 19 to 36 |
| 5 | Long | 18 | 37 to 54 |
| 6 | Longest | 32 | 55 to 86 |
| 7 | Longest | 32 | 87 to 118 |
Read the last column carefully. The atomic numbers run on without a break from one period to the next, which is exactly what the modern periodic law demands — the elements are laid out in one unbroken order of increasing atomic number, and the rows are simply where that order has been cut.
1.2.2.2 Typical Elements Of The Third Period
The third period elements — sodium, magnesium, aluminium, silicon, phosphorus, sulphur, chlorine and argon — summarise the properties of their respective groups, and are often considered typical elements.
This is why so many examples in the rest of the chapter are drawn from the third period. It is short enough to write out in a single line, long enough to run all the way from a reactive metal to a noble gas, and free of the transition elements that interrupt the longer periods.
1.3 Periodicity
Walk along a period and you will find the properties of the elements changing steadily — a soft reactive metal at one end, a gas that reacts with almost nothing at the other. Then start the next period, and the same story begins again.
Nothing about the elements themselves repeats. Their atomic numbers keep climbing, and no element is ever met twice. What repeats is the pattern of behaviour.
The intervals are not random either. Recurrence in properties is seen after a difference of 2, 8, 8, 18, 18 and 32 in atomic numbers, which is exactly the sequence of the period lengths in Table 1.2.
1.3.1 Cause Of Periodicity
The cause of periodicity is the recurrence of similar electronic configuration — that is, elements having the same number of electrons in the outermost orbit, the valence electrons.
Follow one group down the table and you can watch this happening. In a particular group, the electrons in the outermost orbit remain the same, so the electronic configuration is similar even though the atoms grow larger and heavier.
Since the chemical properties of elements depend upon the number of electrons in their outermost shell, elements of the same group have similar properties. This is the whole reason a column of the table is worth drawing.
Diagram
Fig. 1.3-1 — Why Similar Elements Recur At Regular Intervals A horizontal number line of atomic numbers from 1 to 20, each number in its own small box with the element symbol beneath it. Shade the boxes of lithium (3), sodium (11) and potassium (19) in one colour and tag them ‘1 valence electron’. Shade the boxes of helium (2), neon (10) and argon (18) in a second colour and tag them ‘outermost shell complete’. Draw curved arrows above the line joining 3 to 11 and 11 to 19, each labelled ‘+8’, and similar arrows below the line joining 2 to 10 and 10 to 18, each labelled ‘+8’. Caption strip along the base: ‘the elements never repeat — the electron arrangement does’.
1.4 Shells (Orbits) And Valency
Two numbers fix everything about an element’s place in the table: how many shells its atoms have, and how many electrons sit in the outermost one. This section follows both of them as you move down a group and across a period.
Electrons revolve around the nucleus in certain definite circular paths. These paths are called orbits, or shells, and they are counted outwards from the nucleus as K, L, M, N and so on.
1.4.1 Number Of Shells
1.4.1.1 Down A Group
Moving down a group, the number of shells increases successively — one by one — such that the number of shells that an element has equals the number of the period to which that element belongs.
The halogens show this clearly, as set out in Table 1.3. Every one of them keeps seven electrons in its outermost shell; what changes is the number of shells holding them.
Table 1.3 — Number of shells in the elements of group 17
| Element (atomic number) | Number of shells | Electronic configuration (K L M N O P) | Period to which the element belongs |
|---|---|---|---|
| F (9) | 2 | 2, 7 | Second |
| Cl (17) | 3 | 2, 8, 7 | Third |
| Br (35) | 4 | 2, 8, 18, 7 | Fourth |
| I (53) | 5 | 2, 8, 18, 18, 7 | Fifth |
| At (85) | 6 | 2, 8, 18, 32, 18, 7 | Sixth |
Read the last two columns together. Fluorine has two shells and sits in period 2; astatine has six shells and sits in period 6. The period number is simply the shell count written at the edge of the row.
1.4.1.2 Across A Period
On moving from left to right in a given period, the number of shells remains the same.
For example, in the second period the number of shells remains two, equal to the number of the period. Similarly, in the third period the number of shells remains three, and so on.
This matters more than it may seem. Because no new shell is added across a period, every extra electron goes into a shell that is already there, while the nucleus quietly gains another proton. Almost every trend in section 1.5 comes out of that one fact.
1.4.2 Valency
When two elements combine, each atom uses only its valence electrons — giving them away, taking others in, or sharing them. The number of electrons an atom handles in this way is its combining capacity.
Why should an atom bother? Because a completely filled outermost shell is a stable arrangement, as the noble gases show. An atom reacts in whichever way brings it to that arrangement by the shortest route.
1.4.2.1 Valency Down A Group
On moving down a given group, the number of electrons in the outermost shell — the valence electrons — remains the same. Therefore valency, in a group, also remains the same.
Every halogen in Table 1.3 has seven valence electrons and needs one more, so every halogen has a valency of one, from fluorine at the top to astatine at the bottom.
1.4.2.2 Valency Across A Period
In a given period, the number of electrons in the valence shell increases from left to right. Valency, however, does not simply increase with it.
The valency increases only upto group 14, where it becomes 4, and then it decreases, so that in group 17 it becomes 1. Table 1.4 shows the rise and the fall across the second period.
Table 1.4 — Valency across the second period
| Groups → | IA (1) | IIA (2) | IIIA (13) | IVA (14) | VA (15) | VIA (16) | VIIA (17) | Zero (18) |
|---|---|---|---|---|---|---|---|---|
| Element of the 2nd period | Li | Be | B | C | N | O | F | Ne |
| Atomic number | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 |
| Electronic configuration (K L) | 2, 1 | 2, 2 | 2, 3 | 2, 4 | 2, 5 | 2, 6 | 2, 7 | 2, 8 |
| Number of shells | 2 | 2 | 2 | 2 | 2 | 2 | 2 | 2 |
| Valency | 1 | 2 | 3 | 4 | 3 | 2 | 1 | 0 |
| Formula of hydride | \(\mathrm{LiH}\) | \(\mathrm{BeH_2}\) | \(\mathrm{BH_3}\) | \(\mathrm{CH_4}\) | \(\mathrm{NH_3}\) | \(\mathrm{H_2O}\) | \(\mathrm{HF}\) | — |
Plotted as a graph, the valency row of Table 1.4 makes a clear peak, and Fig. 1.4-1 shows its shape.
Diagram
Fig. 1.4-1 — The Rise And Fall Of Valency Across A Period Bar chart. Horizontal axis: the second-period elements in order Li, Be, B, C, N, O, F, Ne, with the number of valence electrons written under each symbol as 1, 2, 3, 4, 5, 6, 7, 8. Vertical axis: valency, scale 0 to 4. Bar heights 1, 2, 3, 4, 3, 2, 1, 0. Shade the first four bars in one colour with a bracket labelled ‘electrons donated’, and the next three in a second colour with a bracket labelled ‘electrons accepted’. Mark the Ne bar of zero height with the tag ‘octet already complete’. Draw a dashed outline over the tops of the bars to bring out the peak at carbon.
The reason for the fall lies in the shortest route to a full shell. If the number of electrons present in the outermost shell is 1, 2, 3 or 4, the atom finds it easier to donate them, and its valency is 1, 2, 3 or 4 respectively.
If the number of electrons present in the outermost shell is 5, 6 or 7, donating them all would be far harder than accepting the few that are missing. Such an atom accepts electrons instead, and its valency is \(8-5=3\), \(8-6=2\) and \(8-7=1\) respectively.
Valency is a combining capacity — a count of electrons handled — so it is always positive. It never carries a plus or a minus sign, whichever way the electrons actually move.
1.5 Periodic Properties
The properties of the elements that are directly or indirectly related to their electronic configurations, and that show a regular gradation as we move across a period or down a group, are the periodic properties of the elements.
Six of them are studied at this level:
- atomic size (atomic radius),
- metallic character,
- non-metallic character,
- ionisation potential (ionisation energy),
- electron affinity,
- electronegativity.
They are not six unrelated facts. Every one of them is decided by a tug-of-war between two things you already know from section 1.4 — the number of shells holding the valence electrons out at a distance, and the charge on the nucleus pulling them in. Once you can weigh those two against each other, all six trends follow.
1.5.1 Atomic Size (Atomic Radius)
An atom has no hard surface, so its size has to be defined by where its outermost electrons are.
Atomic radius can also be defined as half the inter-nuclear distance between the combined atoms in a molecule. Fig. 1.5-1 shows both ways of measuring it.
Diagram
Fig. 1.5-1 — Two Ways Of Measuring Atomic Radius Two panels side by side. Left panel: a single atom drawn as two concentric circles with a shaded nucleus at the centre; a straight arrow runs from the centre of the nucleus to the outer circle, labelled ‘atomic radius’; label the shaded centre ‘nucleus’. Right panel: two identical atoms of the same element drawn as touching, slightly overlapping circles, each with a dot at its centre; a horizontal double-headed arrow joins the two centre dots and is labelled ‘inter-nuclear distance’; a bracket over one half of that arrow is labelled ‘atomic radius = half the inter-nuclear distance’; label each circle ‘atom’.
Atoms are far too small for metres to be a convenient unit, so two smaller units are used:
\[1\ \text{\AA}\ (\text{angstrom}) = 10^{-10}\ \text{m}\]
\[1\ \text{pm}\ (\text{picometre}) = 10^{-12}\ \text{m}\]
A picometre is a hundredth of an angstrom, so an atomic radius of \(1.52\ \text{\AA}\) is the same as \(152\ \text{pm}\).
1.5.1.1 What Atomic Size Depends On
Atomic size depends upon two things: the number of shells, and the nuclear charge.
An increase in the number of shells increases the size of an atom, because the distance between the outermost shell and the nucleus increases. Each new shell is laid outside the last one.
The second factor pulls the other way. Nuclear charge is the positive charge present in the nucleus of an atom, which is equal to the number of protons in the nucleus — that is, to the atomic number.
An increase in nuclear charge decreases the size of the atom, because the electrons are then attracted towards the nucleus with a greater force, which brings the outermost shell closer to the nucleus.
1.5.1.2 Trends In Atomic Size
Down a group. In a group, the size of an atom increases as one proceeds from top to bottom. This is due to the successive addition of shells, which overweighs the increased nuclear charge, as one moves from one period to the next in a group.
Table 1.5 shows this for group 1, where hydrogen has the smallest size of all.
Table 1.5 — Atomic radii of the elements of group 1
| Element | Atomic number | Atomic radius | Electronic configuration (K L M N O P) |
|---|---|---|---|
| Hydrogen | 1 | 37 pm | 1 |
| Lithium | 3 | 152 pm | 2, 1 |
| Sodium | 11 | 186 pm | 2, 8, 1 |
| Potassium | 19 | 231 pm | 2, 8, 8, 1 |
| Rubidium | 37 | 244 pm | 2, 8, 18, 8, 1 |
| Caesium | 55 | 262 pm | 2, 8, 18, 18, 8, 1 |
The halogens behave in the same way, with fluorine the smallest of them:
\[\mathrm{F}\ (64\ \text{pm}) < \mathrm{Cl}\ (99\ \text{pm}) < \mathrm{Br}\ (114\ \text{pm}) < \mathrm{I}\ (133\ \text{pm}) < \mathrm{At}\ (140\ \text{pm})\]
Across a period. In a period, the size of an atom decreases from left to right. This is because the nuclear charge, or nuclear pull, increases due to an increase in atomic number from left to right in the same period, which brings the outermost shell closer to the nucleus. Meanwhile the number of electron shells in a period remains the same, so there is nothing to push the outer electrons away.
Therefore, in the first period, hydrogen has a larger size than helium. In the second period, lithium has the largest atomic size while fluorine has the smallest. In the third period, the sodium atom is the largest in size and the chlorine atom is the smallest:
\[\mathrm{Na} > \mathrm{Mg} > \mathrm{Al} > \mathrm{Si} > \mathrm{P} > \mathrm{S} > \mathrm{Cl}\]
\[186\ \text{pm} \quad 160\ \text{pm} \quad 143\ \text{pm} \quad 117\ \text{pm} \quad 110\ \text{pm} \quad 104\ \text{pm} \quad 99\ \text{pm}\]
Table 1.6 gives the same fall for the second period, and Fig. 1.5-2 draws it.
Table 1.6 — Atomic radii across the second period
| Element | Lithium | Beryllium | Boron | Carbon | Nitrogen | Oxygen | Fluorine | Neon |
|---|---|---|---|---|---|---|---|---|
| Atomic number | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 |
| Atomic radius | 152 pm | 112 pm | 88 pm | 77 pm | 70 pm | 66 pm | 64 pm | 112 pm |
| Electronic configuration (K L) | 2, 1 | 2, 2 | 2, 3 | 2, 4 | 2, 5 | 2, 6 | 2, 7 | 2, 8 |
Diagram
Fig. 1.5-2 — How Atomic Size Changes Down A Group And Across A Period Two sets of shell diagrams, all drawn to one common scale so the sizes can be compared by eye. Left of the panel, a vertical column running downwards: lithium, sodium, potassium, rubidium and caesium, each with the correct number of concentric shells and a small shaded nucleus, the circles growing steadily larger; a long downward arrow beside them labelled ‘down group 1: size increases — a new shell each time’. Along the base, a horizontal row running left to right: lithium, beryllium, boron, carbon, nitrogen, oxygen, fluorine, each with two shells only, the circles shrinking steadily; a long rightward arrow beneath labelled ‘across period 2: size decreases — same shells, stronger nuclear pull’. Write the atomic radius under each atom.
Neon in Table 1.6 breaks the fall, and the noble gases do this in every period. As an exception, the size of the atoms of the inert gases is bigger than that of the halogens of the same period. This is because the outermost shell of an inert gas is complete: it holds the maximum number of electrons, so the electronic repulsions between them are maximum, and the effect of nuclear pull over the valence electrons is not seen.
Hence the atom of an inert gas is bigger than the halogen beside it, but it is not bigger than the group 1 element of the same period. Helium, with a single filled shell, has the smallest size in the periodic table.
1.5.1.3 Sizes Of Cations, Anions And Isoelectronic Ions
An atom changes size the moment it becomes an ion, and it changes in a direction you can predict.
A cation is always smaller than the parent atom from which it is formed. A cation is formed by the loss of electron or electrons, so the protons now outnumber the electrons. The remaining electrons are strongly attracted by the nucleus and are pulled inward, and the size decreases:
\[\mathrm{Na} - e^{-} \rightarrow \mathrm{Na^{+}}\]
Sodium is 2, 8, 1 and the sodium ion is 2, 8, so an entire shell has also been emptied. The sodium ion is smaller than the sodium atom.
An anion is larger than the parent atom. An anion is formed by the gain of electron or electrons, so the electrons now outnumber the protons. These electrons repel each other, which reduces the overall attraction between the nucleus and the electrons. The effective nuclear charge is reduced and the size becomes larger:
\[\mathrm{Cl} + e^{-} \rightarrow \mathrm{Cl^{-}}\]
Chlorine is 2, 8, 7 and the chloride ion is 2, 8, 8. Here no shell has been added, yet the chloride ion is bigger than the chlorine atom, purely because of the added repulsion. Fig. 1.5-3 shows both changes together.
Diagram
Fig. 1.5-3 — Why A Cation Shrinks And An Anion Swells Two rows. Top row: sodium atom (three shells, 2-8-1, nucleus marked 11 p, 11 e) with an arrow labelled ‘loses 1 electron’ leading to a distinctly smaller sodium ion drawn with two shells (2-8, nucleus 11 p, 10 e); caption strip under the row ‘more protons than electrons → electrons pulled in → smaller’. Bottom row: chlorine atom (three shells, 2-8-7, nucleus 17 p, 17 e) with an arrow labelled ‘gains 1 electron’ leading to a slightly larger chloride ion (2-8-8, nucleus 17 p, 18 e), with small repulsion arrows drawn between the electrons of the outer shell; caption strip ‘more electrons than protons → extra repulsion → larger’. Draw both members of each pair to scale so the size change is obvious.
Isoelectronic ions are ions having the same number of electrons. Since the electron count is fixed, their sizes depend only upon the nuclear charge, that is, the number of protons. The greater the nuclear charge, the greater the nuclear pull, and the smaller the ion.
Table 1.7 — Sizes of some isoelectronic ions
| Isoelectronic ion | \(\mathrm{Mg^{2+}}\) | \(\mathrm{Na^{+}}\) | \(\mathrm{F^{-}}\) | \(\mathrm{O^{2-}}\) |
|---|---|---|---|---|
| Number of electrons | 10 | 10 | 10 | 10 |
| Number of protons | 12 | 11 | 9 | 8 |
| Size in Å | 0.65 | 0.95 | 1.26 | 1.42 |
All four ions have the electron arrangement 2, 8. Reading from left to right the proton count falls from 12 to 8, and the size climbs steadily from \(0.65\ \text{\AA}\) to \(1.42\ \text{\AA}\). Ten electrons held by twelve protons are gripped far more tightly than the same ten held by eight.
1.5.2 Metallic Character
Those elements which have a tendency to lose their valence electrons — the electrons of the outermost orbit — and form a positive ion are considered metals. Metallic character is a measure of how readily an element does this.
\[\mathrm{Na} - e^{-} \rightarrow \mathrm{Na^{+}} \qquad (2,\,8,\,1) \rightarrow (2,\,8)\]
\[\mathrm{Mg} - 2e^{-} \rightarrow \mathrm{Mg^{2+}} \qquad (2,\,8,\,2) \rightarrow (2,\,8)\]
In each case the atom ends up with a complete outermost shell, which is the arrangement it was reaching for.
Elements which lose electrons to complete their octet are reducing agents, since the electrons they give away are taken up by something else. Metals are good reducing agents, and the greater the tendency to lose electrons, the stronger the reducing agent is.
1.5.2.1 What Metallic Character Depends On
The metallic character of elements depends on the same two quantities as size.
Atomic size. The greater the atomic size, the farther is the outermost orbit, and thus the lesser is the nuclear pull exerted on it. As a result, electrons can be removed more easily from the valence shell, making the element more metallic.
Nuclear charge. The greater the nuclear charge, the greater is the force exerted by the nucleus on the electrons of the outermost orbit. This makes it difficult to remove those electrons, and thus metallic nature decreases.
1.5.2.2 Trends In Metallic Character
Down a group. On moving down a group, the atomic size increases, and the nuclear charge also increases. The effect of the increased atomic size is greater as compared to the increased nuclear charge. Therefore the tendency to lose electrons increases, and elements can lose electrons easily. Thus metallic nature increases as one moves down a group.
Elements at the bottom of a group are the most metallic. The atomic size is large, so the electrons are loosely held and ions are readily formed, which also makes these elements more reactive. In group 1, lithium is the least metallic while caesium is the most metallic element. Francium, below caesium, is a radioactive element, so its properties are not known.
Across a period. On moving across a period, the nuclear pull increases due to the increase in atomic number, and thus the atomic size decreases. Hence elements cannot lose electrons easily. Therefore the metallic nature decreases across a period, moving from left to right, due to the increasing nuclear charge.
In the second period, lithium is the most metallic element, and the change from metal to non-metal can be followed element by element:
\[\underbrace{\mathrm{Li}\quad \mathrm{Be}}_{\text{metals}} \quad \underbrace{\mathrm{B}}_{\text{metalloid}} \quad \underbrace{\mathrm{C}\quad \mathrm{N}\quad \mathrm{O}\quad \mathrm{F}}_{\text{non-metals}}\]
An element such as boron, which shows some properties of metals and some of non-metals, is called a metalloid. In the third period, sodium is the most metallic element, and Table 1.8 follows the same change across that period.
Table 1.8 — Metallic property and valency across the third period
| Group | 1 | 2 | 13 | 14 | 15 | 16 | 17 | 18 |
|---|---|---|---|---|---|---|---|---|
| Element | Na | Mg | Al | Si | P | S | Cl | Ar |
| Atomic number | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 |
| Electronic configuration (K L M) | 2, 8, 1 | 2, 8, 2 | 2, 8, 3 | 2, 8, 4 | 2, 8, 5 | 2, 8, 6 | 2, 8, 7 | 2, 8, 8 |
| Metallic property | Metal | Metal | Metal | Metalloid | Non-metal | Non-metal | Non-metal | Noble gas |
| Valency | 1 | 2 | 3 | 4 | 3 | 2 | 1 | 0 |
Hydrogen deserves a note of its own here. It is a non-metal, and its atom has no neutron — just one proton and one electron. On losing that electron it forms a cation containing nothing but a single proton, so the hydrogen ion can also be called a proton:
\[\mathrm{H} - e^{-} \rightarrow \mathrm{H^{+}}\]
Since hydrogen can also lose its one electron, it is a good reducing agent too.
1.5.3 Non-Metallic Character
Those elements which have a tendency to gain electrons, in order to attain an octet in their outermost orbit, are considered as non-metals. Non-metallic character measures how strongly an element does this.
Non-metals usually have 5, 6 or 7 electrons in their outermost orbits. They can gain 3, 2 or 1 electron respectively, in order to attain the octet and complete their valence shell, and so form an anion.
\[\mathrm{Cl} + e^{-} \rightarrow \mathrm{Cl^{-}} \qquad (2,\,8,\,7) \rightarrow (2,\,8,\,8)\]
\[\mathrm{O} + 2e^{-} \rightarrow \mathrm{O^{2-}} \qquad (2,\,6) \rightarrow (2,\,8)\]
An element that takes electrons in this way is taking them from something else, which is thereby oxidised. Non-metals are therefore oxidising agents, which is the exact mirror of the statement that metals are reducing agents.
1.5.3.1 What Non-Metallic Character Depends On
Atomic size. The smaller the atomic size, the greater is the nuclear pull and the greater is the tendency to gain electrons. Therefore the element is more non-metallic by nature.
Nuclear charge. The greater the nuclear charge, the greater is the tendency to gain electrons, and hence the more non-metallic is the element.
1.5.3.2 Trends In Non-Metallic Character
Down a group. The atomic size increases due to the addition of new shells over successive periods. Though the nuclear charge increases as well, because of the rise in atomic number, the effect of the increasing atomic size is greater. Therefore non-metallic nature decreases down the group.
Group 14 shows the change completely, running from a non-metal at the top to metals at the bottom:
- C — non-metal
- Si — metalloid
- Ge — metalloid
- Sn — metal
- Pb — metal
Across a period. On moving across a period, the tendency to gain electrons increases due to an increase in nuclear pull and a decrease in the atomic size. Therefore non-metallic character increases across a period, from left to right. For example, in the third period:
\[\mathrm{Na} < \mathrm{Mg} < \mathrm{Al} < \mathrm{Si} < \mathrm{P} < \mathrm{S} < \mathrm{Cl}\]
1.5.3.3 Periodicity In The Nature Of Oxides
The nature of the oxides also shows periodicity, and it is a useful check on the metallic or non-metallic character of the element itself.
Across a period, the oxides of the elements show a decreasing basic nature and finally become acidic. Table 1.9 sets out the oxides of the third period.
Table 1.9 — Nature of the oxides of the third period
| Oxide | \(\mathrm{Na_2O}\) | \(\mathrm{MgO}\) | \(\mathrm{Al_2O_3}\) | \(\mathrm{SiO_2}\) | \(\mathrm{P_2O_5}\) | \(\mathrm{SO_3}\) | \(\mathrm{Cl_2O_7}\) |
|---|---|---|---|---|---|---|---|
| Nature | Strongly basic | Basic | Amphoteric | Feebly acidic | Acidic | More acidic | Most acidic |
Aluminium oxide sits at the turning point, and is described as amphoteric — it behaves as a basic oxide towards an acid and as an acidic oxide towards an alkali. Down a group, the basic nature of the oxides of metals increases, which is exactly what you would expect from metallic character increasing downwards.
1.5.4 Chemical Reactivity
The reactivity of elements depends upon their tendency to lose or gain electrons to complete their outermost orbit. The greater the tendency to lose electrons, the greater is the reactivity in the case of metals. Similarly, the greater the tendency to gain electrons, the greater is the reactivity of non-metals.
1.5.4.1 Reactivity Across A Period
On moving from left to right in a period, the chemical reactivity of elements first decreases and then increases.
Follow the third period to see why. The group 1 element can lose electrons easily as compared to group 2. Group 2 elements can also lose electrons easily, but only in comparison to group 13, and not at all in comparison to group 1. As the tendency to lose electrons decreases, reactivity decreases, and thus silicon is the least reactive element in the third period.
From that point the other tendency takes over. As we move from phosphorus to chlorine, the tendency to gain electrons increases, and hence reactivity increases again. The result is the V-shaped pattern drawn in Fig. 1.5-4.
Diagram
Fig. 1.5-4 — Reactivity Across The Third Period A horizontal strip showing the third-period elements in order: Na, Mg, Al, Si, P, S, Cl. Beneath the strip draw a broad V-shaped curve whose height represents reactivity — high above Na, falling to its lowest point above Si, and rising again to a high point above Cl. Label the left arm ‘tendency to lose electrons decreases → reactivity decreases’; label the right arm ‘tendency to gain electrons increases → reactivity increases’. Tag Na ‘most reactive metal’, Si ‘least reactive’, Cl ‘most reactive non-metal’.
1.5.4.2 Reactivity Down A Group
The tendency to lose electrons increases down a group. Since chemical reactivity in metals depends upon the tendency to lose electrons, reactivity increases on going down the group, and the most reactive metal is at the bottom of group 1.
The chemical reactivity of non-metals decreases on going down the group, as it depends upon the tendency to gain electrons, which decreases downwards. The most reactive non-metal is therefore at the top of group 17, that is, fluorine.
1.5.5 Gradation In Physical Properties
Melting point, boiling point and density are not decided by the valence electrons alone, so their trends are less tidy than the ones above. They still show a clear gradation.
The melting and boiling points of metals decrease on going down the group.
Table 1.10 — Melting and boiling points of some alkali metals
| Metal | Melting point | Boiling point |
|---|---|---|
| Li | 180.5 °C | 1347 °C |
| Na | 94.5 °C | 883 °C |
| K | 63.5 °C | 774 °C |
The melting and boiling points of non-metals increase on going down the group, which is why the halogens change state as you descend group 17.
Table 1.11 — Melting points, boiling points and states of the halogens
| Non-metal | Melting point | Boiling point | Physical state |
|---|---|---|---|
| Fluorine | −219.6 °C | −187 °C | gas |
| Chlorine | −101 °C | −34.6 °C | gas |
| Bromine | −7.2 °C | +58.8 °C | liquid |
| Iodine | +113.6 °C | +183 °C | solid |
Across a period, from left to right, melting point and boiling point usually increase upto group 14 and then decrease. Density across a period increases gradually to a maximum, and then a slight decrease may be noticed. Table 1.12 shows both across the third period.
Table 1.12 — Melting point and density across the third period
| Element of 3rd period | Na | Mg | Al | Si | P | S |
|---|---|---|---|---|---|---|
| Melting point (°C) | 98 | 650 | 660 | 1410 | 44.2 | 115.2 |
| Density (g/cc) | 1.0 | 1.7 | 2.7 | 2.3 | 1.8 | 2.1 |
Down a group, the density of elements increases gradually.
Table 1.13 — Density down group 1
| Element | Li | Na | K | Rb | Cs |
|---|---|---|---|---|---|
| Density (g/cc) | 0.54 | 0.97 | 0.86 | 1.53 | 1.87 |
Notice that potassium’s density is a little lower than sodium’s. Measured values do not always fall in a perfect line, and a good trend is one that holds overall, not one that has no exception anywhere.
1.5.6 Ionisation Potential
The electrons in an atom are attracted by the positively charged nucleus. So if we want to remove an electron from an atom, some energy has to be supplied to overcome this strong attraction. Once the electron has been removed, the atom acquires one unit of positive charge and becomes a positive ion, or cation.
\[\mathrm{M}\text{(g)} + \text{I.E.} \rightarrow \mathrm{M^{+}}\text{(g)} + e^{-}\]
Here \(\mathrm{M}\text{(g)}\) represents an isolated gaseous atom and \(\mathrm{M^{+}}\text{(g)}\) represents the gaseous cation of the same atom. Every word of that definition is doing work: gaseous and isolated, because a neighbouring atom would interfere with the measurement, and most loosely bound, because it is the outermost electron that leaves first.
Ionisation energy is measured in electron volts per atom (eV/atom), and its S.I. unit is kilojoule per mole (\(\text{kJ mol}^{-1}\)).
1.5.6.1 What Ionisation Energy Depends On
Atomic size. The greater the atomic size, the lesser is the force of attraction. Since the electrons of the outermost shell lie further away from the nucleus, their removal becomes easier, and the ionisation energy required is less.
Nuclear charge. The greater the nuclear charge, the greater is the attraction for the electrons of the outermost shell. Those electrons are then more firmly held, because of which greater energy is required to remove them.
1.5.6.2 Trends In Ionisation Energy
Across a period. The ionisation energy tends to increase as one moves from left to right across a period, with exceptions. The number of protons in the nucleus increases, leading to a stronger positive charge. This increased nuclear charge attracts the electrons more strongly, making it harder to remove an electron, and thus more energy is required.
Table 1.14 — Ionisation energy across the second period
| Element of 2nd period | Li | Be | B | C | N | O | F | Ne |
|---|---|---|---|---|---|---|---|---|
| Ionisation energy (kJ mol⁻¹) | 520 | 899* | 801 | 1086 | 1402* | 1314 | 1681 | 2080 |
Table 1.15 — Ionisation energy across the third period
| Element of 3rd period | Na | Mg | Al | Si | P | S | Cl | Ar |
|---|---|---|---|---|---|---|---|---|
| Ionisation energy (kJ mol⁻¹) | 496 | 737* | 577 | 786 | 1011* | 999 | 1256 | 1520 |
The values marked with an asterisk are the exceptions. Beryllium and nitrogen sit higher than the elements that follow them, and magnesium and phosphorus do the same in the third period; the overall climb from left to right is unmistakable all the same.
Down a group. There is an increase in atomic number, and so in nuclear charge, and an increase in atomic size down the group due to the addition of extra shells. This increase in the atomic size overcomes the effect of the increase in the nuclear charge.
Therefore ionisation energy decreases with an increase in the atomic size — that is, it decreases as one moves down a group.
Table 1.16 — Ionisation energy down group 1
| Element | Li | Na | K | Rb | Cs |
|---|---|---|---|---|---|
| Ionisation energy (kJ mol⁻¹) | 520 | 496 | 419 | 403 | 374 |
Helium has the highest ionisation energy of any element, \(2372.0\ \text{kJ mol}^{-1}\), which fits its tiny size and complete shell. Caesium has the lowest value among the elements in ordinary use; the ionisation energy of francium is not determined correctly, as it is radioactive. As a general rule, metals usually have low ionisation energy whereas non-metals have high ionisation energy.
1.5.7 Electron Affinity
To remove an electron from an atom, energy is required. Something close to the opposite happens when an extra electron is added to an atom: energy is released.
\[\mathrm{X}\text{(g)} + e^{-} \rightarrow \mathrm{X^{-}}\text{(g)} + \text{E.A.}\]
Here \(\mathrm{X}\) is any element taken in its gaseous state. For chlorine:
\[\mathrm{Cl}\text{(g)} + e^{-} \rightarrow \mathrm{Cl^{-}}\text{(g)} + 349\ \text{kJ mol}^{-1}\]
Since that energy leaves the atom rather than being supplied to it, the electron affinity of chlorine is written as \(-349\ \text{kJ mol}^{-1}\). Fig. 1.5-5 sets ionisation energy and electron affinity side by side.
Diagram
Fig. 1.5-5 — Energy In, Energy Out Two panels side by side sharing one caption strip. Left panel: a single gaseous atom labelled ‘M (g)’ with an inward arrow labelled ‘energy supplied = ionisation energy’ striking it, and an electron flying away from the outer shell; the product drawn to the right as a smaller ion labelled ‘M⁺ (g)’. Right panel: a single gaseous atom labelled ‘X (g)’ with an electron arriving from the left along an arrow labelled ‘electron added’, and an outward wavy arrow labelled ‘energy released = electron affinity’; the product drawn to the right as a larger ion labelled ‘X⁻ (g)’. Caption strip: ‘taking an electron away costs energy; taking one in gives energy back’.
1.5.7.1 What Electron Affinity Depends On
Atomic size. The smaller the atomic size, the greater the electron affinity, because the effective attractive force between the nucleus and the valence electrons is greater in smaller atoms, and thus the electron is more readily accepted.
Nuclear charge. The greater the nuclear charge, the greater is the electron affinity, because an increase in nuclear charge increases the effective attractive force on the valence electrons.
1.5.7.2 Trends In Electron Affinity
Across a period. In a period, from left to right, the atomic size decreases and the nuclear charge increases, so the electron affinity increases. Thus electron affinity is the highest for the group 17 elements, the halogens, and the least for the group 1 elements, the alkali metals.
Table 1.17 — Electron affinity values in kJ mol⁻¹
| Period | 1 (IA) | 2 (IIA) | 13 (IIIA) | 14 (IVA) | 15 (VA) | 16 (VIA) | 17 (VIIA) | 18 (Zero) |
|---|---|---|---|---|---|---|---|---|
| I | H −73 | — | — | — | — | — | — | He +48 |
| II | Li −60 | Be* | B −88 | C −122 | N* | O −141 | F −328 | Ne* |
| III | Na −53 | Mg* | Al −50 | Si −119 | P −74 | S −200 | Cl −349 | Ar* |
The inert gases have zero electron affinity or positive values, because their stable electronic configuration makes it difficult for them to accept an electron. Groups 2 and 15, marked with an asterisk above, do not show regular trends in electron affinity values, and are exceptions.
Down a group. Moving from the top to the bottom in a group, the atomic size increases more than the nuclear charge, causing a net decrease in electron affinity: H \(-73\), Li \(-60\), Na \(-53\), K \(-48\), Rb \(-47\ \text{kJ mol}^{-1}\).
There is a well-known exception at the very top of two groups. In group 17, fluorine has lower electron affinity than chlorine, and in group 16, oxygen has lower electron affinity than sulphur. This is because the fluorine and oxygen atoms are very small, so there are strong inter-electronic repulsions inside the crowded outer shell, and the incoming electron does not feel much attraction.
Electron affinity does not show a perfectly regular trend along a period or a group, because of the number of exceptions. In general it becomes less negative in going from top to bottom in a group, and more negative in going from left to right in a period.
1.5.8 Electronegativity
When two different atoms share a pair of electrons, the pair is rarely shared fairly. One of the two atoms pulls it closer.
Electronegativity is a dimensionless property, since it is only a tendency. It carries no unit, because it does not measure an amount of anything; it only indicates the net result of the tendencies of different elements to attract the bond-forming electron pair.
Electronegativity is measured on several scales. The most widely used scale was devised by Linus Pauling in 1932, in which the highest value, for fluorine, is taken as 4.0 and the lowest, for caesium, as 0.7.
Table 1.18 — Electronegativity values on the Pauling scale
| Group → | 1 | 2 | 13 | 14 | 15 | 16 | 17 |
|---|---|---|---|---|---|---|---|
| H 2.1 | — | — | — | — | — | — | |
| Li 1.0 | Be 1.5 | B 2.0 | C 2.5 | N 3.0 | O 3.5 | F 4.0 | |
| Na 0.9 | Mg 1.2 | Al 1.5 | Si 1.8 | P 2.1 | S 2.5 | Cl 3.0 | |
| K 0.8 | Ca 1.0 | Ga 1.6 | Ge 1.8 | As 2.0 | Se 2.4 | Br 2.8 | |
| Rb 0.8 | Sr 1.0 | In 1.7 | Sn 1.8 | Sb 1.9 | Te 2.1 | I 2.5 | |
| Cs 0.7 | Ba 0.9 | Tl 1.8 | Pb 1.9 | Bi 1.9 | Po 2.0 | At 2.2 |
Read Table 1.18 across and the values rise; read it down and they fall. Those two readings are the whole of what follows.
1.5.8.1 What Electronegativity Depends On
Atomic size. The greater the size of the atom, the lesser the electronegativity, since the electrons, being farther away from the nucleus, experience a lesser force of attraction.
Nuclear charge. The greater the nuclear charge, the greater the electronegativity, because an increase in nuclear charge causes electron attraction with a greater force.
1.5.8.2 Trends In Electronegativity
Across a period. Since the nuclear charge increases due to an increase in atomic number, electronegativity increases from left to right in a period. For example, in the second period, electronegativity increases from lithium at 1.0 to fluorine at 4.0.
Down a group. There is an increase in atomic number down a group, so the nuclear charge increases; but due to the addition of extra shells, the atomic size increases as well. The effect of the increase in the atomic size overcomes the effect of the increase in the nuclear charge, hence electronegativity decreases down a group. In the first group, for example, it decreases significantly from lithium downwards to francium.
The noble gases have a complete octet, so they do not have the tendency to attract electrons.
Reading The Six Trends Together
The last four properties are really four ways of asking one question: how tightly does this atom hold electrons, and how badly does it want more?
The greater the value of the electron affinity, the easier it is for an element to gain electrons, and so the more non-metallic, the more electronegative and the more oxidising the element is. Thus oxidising power increases from left to right in a period, and decreases from top to bottom in a group.
Metals lose electrons. They have less ionisation energy, less electron affinity and less electronegativity, and they have large atomic radii as compared to non-metals. Generally, metals show lower electronegativity than non-metals, so metals are electropositive and non-metals are electronegative, and metals are good reducing agents.
One consequence is worth carrying forward. The elements of the second period differ in properties from the other members of their respective groups, and this is due to the small size of their atoms and their high electronegativity.
1.6 Diagonal Relationship And Bridge Elements
Look back at the trends of section 1.5 and notice something about their directions. Properties which have an increasing trend in a group generally have a decreasing trend in the period, and the other way round.
Now take one step to the right and one step down. The step to the right raises the nuclear pull and shrinks the atom; the step down adds a shell and enlarges it. The two changes work against each other, and to a useful extent they cancel.
That is why the elements of the second period show a resemblance in properties with the elements of the next group of the third period, present at the right diagonal of the element. This leads to a diagonal relationship, as shown in Fig. 1.6-1.
Diagram
Fig. 1.6-1 — The Diagonal Relationship A small grid, four columns by two rows. Column headings across the top: ‘Group → 1, 2, 13, 14’. Row label ‘Period 2’ on the left of the upper row, holding the symbols Li, Be, B, C in order. Row label ‘Period 3’ on the left of the lower row, holding Na, Mg, Al, Si in order. Draw three bold diagonal arrows sloping down and to the right: Li to Mg, Be to Al, and B to Si. Tag the pairs ‘similar properties’. Add a small key beneath: ‘one step right = smaller atom, stronger pull; one step down = larger atom, weaker pull; the two roughly cancel along the diagonal’. Label the C box at the end of the upper row ‘bridge element’.
The pairs joined by these arrows are lithium and magnesium, beryllium and aluminium, and boron and silicon. Such elements are called bridge elements, because each one bridges its own group and the group to its right.
1.7 Atomic Number And Mass Number
Two numbers are attached to every atom, and it is worth being clear about what each one is counting before you use either.
1.7.1 Atomic Number (Z)
The atomic number of an element, defined in section 1.1.2, is equal to the number of protons in the nucleus. It is the unique property of an element, because no two elements have the same atomic number.
\[\text{Atomic number}\ (Z) = \text{number of protons} = \text{number of electrons}\]
The second equality holds only for a neutral atom, and it is the one that makes the atomic number so useful. Knowing \(Z\), you can do three things at once.
- It gives the electronic configuration of an element. An element with atomic number 13 will have the configuration 2, 8, 3.
- It helps in finding the position of an element in the periodic table.
- It identifies the element beyond any doubt, whatever its mass turns out to be.
For example, the element with atomic number 17 will have the electronic configuration 2, 8, 7. This element will be placed in the third period of group 17, because it has three energy shells and seven electrons in its outermost shell.
1.7.2 Mass Number (A)
Almost all of an atom’s mass sits in its nucleus, in the protons and the neutrons together. The electrons are far too light to count.
\[\text{Mass number}\ (A) = \text{number of protons}\ (p) + \text{number of neutrons}\ (n)\]
Rearranged, the same relation gives you the neutron count, which is not measured directly:
\[n = A - Z\]
Both numbers are written with the symbol of the element, as shown in Fig. 1.7-1.
Diagram
Fig. 1.7-1 — How An Atom Is Labelled Left: the symbol ³⁵₁₇Cl drawn large, with the mass number as the upper left index and the atomic number as the lower left index. Draw a leader line from the upper index to a label ‘mass number A = protons + neutrons = 35’ and from the lower index to a label ‘atomic number Z = protons = 17’. Right: a simple sketch of the same atom — a nucleus containing 17 shaded circles marked ‘p’ and 18 plain circles marked ‘n’, surrounded by three shells carrying 2, 8 and 7 electrons; label the nucleus ‘A − Z = 35 − 17 = 18 neutrons’.
A Pattern In The Lighter Elements
A close look at the lighter elements reveals a rough arithmetic pattern between the two numbers.
Elements which have an even number of protons — for example \(^{4}_{2}\mathrm{He}\) and \(^{12}_{6}\mathrm{C}\) — have their mass numbers twice the atomic numbers, except \(^{9}_{4}\mathrm{Be}\) and \(^{40}_{18}\mathrm{Ar}\).
Elements which have an odd number of protons, like \(^{7}_{3}\mathrm{Li}\), \(^{11}_{5}\mathrm{B}\), \(^{19}_{9}\mathrm{F}\) and \(^{23}_{11}\mathrm{Na}\), have their mass numbers twice the atomic numbers plus one, that is \(A = 2Z + 1\), except \(^{14}_{7}\mathrm{N}\) and \(^{1}_{1}\mathrm{H}\).
Table 1.20 sets out the counts for the first twenty elements, and you can check both patterns against it.
Table 1.20 — Arrangement of electrons, protons and neutrons of the elements from atomic number 1 to 20
| Element | Symbol | Atomic number (Z), number of protons or electrons | Mass number (A), protons + neutrons | Number of neutrons (A − Z) |
|---|---|---|---|---|
| Hydrogen | H | 1 | 1 | — |
| Helium | He | 2 | 4 | 2 |
| Lithium | Li | 3 | 7 | 4 |
| Beryllium | Be | 4 | 9 | 5 |
| Boron | B | 5 | 11 | 6 |
| Carbon | C | 6 | 12 | 6 |
| Nitrogen | N | 7 | 14 | 7 |
| Oxygen | O | 8 | 16 | 8 |
| Fluorine | F | 9 | 19 | 10 |
| Neon | Ne | 10 | 20 | 10 |
| Sodium | Na | 11 | 23 | 12 |
| Magnesium | Mg | 12 | 24 | 12 |
| Aluminium | Al | 13 | 27 | 14 |
| Silicon | Si | 14 | 28 | 14 |
| Phosphorus | P | 15 | 31 | 16 |
| Sulphur | S | 16 | 32 | 16 |
| Chlorine | Cl | 17 | 35, 37 | 18, 20 |
| Argon | Ar | 18 | 40 | 22 |
| Potassium | K | 19 | 39 | 20 |
| Calcium | Ca | 20 | 40 | 20 |
Chlorine is given two entries in the last two columns because chlorine atoms come in two kinds, with 18 and 20 neutrons. The proton count is 17 in both, so both are chlorine — which is exactly the point that defeated Mendeleev’s ordering by weight.
1.7.3 Neutron To Proton Ratio And Stability
Read down the last two columns of Table 1.20 and you will see the neutron count creeping ahead of the proton count as the elements get heavier. The ratio of the two decides whether a nucleus holds together.
Elements with an \(n/p\) ratio around 1 are stable — the light metals such as sodium, potassium and calcium are examples.
Elements with an \(n/p\) ratio of 1.5 and above are radioactive, that is, they emit radiations. Such elements are unstable, and the heavy metals like uranium belong here.
1.8 Comparison Of Alkali Metals And Halogens
The two families at the edges of the table make the clearest possible pair to compare. Group 1 sits at the far left with a single electron to give away; group 17 sits at the far right needing a single electron to complete its octet.
Everything else about them follows from that one difference of one electron, as Fig. 1.8-1 shows.
Diagram
Fig. 1.8-1 — One Electron Spare, One Electron Short Left: outline grid of the periodic table with column 1 shaded and labelled ‘alkali metals — 1 valence electron’ and column 17 shaded in a contrasting colour and labelled ‘halogens — 7 valence electrons’; hydrogen’s box left unshaded with the tag ‘not an alkali metal’. Right: two shell diagrams one above the other. Upper — a sodium atom (2, 8, 1) with an arrow showing its single outer electron leaving, product labelled Na⁺, caption ‘loses one → cation, a reducing agent’. Lower — a chlorine atom (2, 8, 7) with an arrow showing one electron arriving into the outer shell, product labelled Cl⁻, caption ‘gains one → anion, an oxidising agent’. Draw a curved arrow from the departing electron of the upper diagram to the arriving electron of the lower diagram, labelled ‘and this is why they combine’.
A word on handling before the comparison itself. Both families are reactive enough to be dangerous. The alkali metals are stored under oil because they attack moist air, and are handled with tongs and cut behind a screen. The halogens are poisonous gases or vapours, and any demonstration involving them is carried out in a fume cupboard. Neither family is material for an open bench.
Table 1.21 compares the two, property by property.
The Reaction With Water, Written Out
The row on water and acids is worth writing as equations, since these are reactions you will use again. Taking \(\mathrm{M}\) to stand for any alkali metal:
alkali metal + water → alkali metal hydroxide + hydrogen
\[2\,\mathrm{M}\text{(s)} + 2\,\mathrm{H_2O}\text{(l)} \rightarrow 2\,\mathrm{MOH}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
alkali metal + dilute hydrochloric acid → alkali metal chloride + hydrogen
\[2\,\mathrm{M}\text{(s)} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow 2\,\mathrm{MCl}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
The hydroxide formed in the first reaction is the strong alkali that gave the family its name, back in section 1.2.1.1. And because reactivity increases down the group, the same reaction that is brisk with lithium is violent with caesium.
The compounds formed in the last row of Table 1.21 are of two kinds. An electrovalent compound is one formed by the complete transfer of electrons from one atom to another, which is what happens when a metal that must lose an electron meets a non-metal that must gain one. Where two non-metals meet, neither can take outright, so the electrons are shared instead and a covalent compound results.
1.9 Summary
- Elements are pure substances made of one kind of atom. Because they are so many, they must be classified before they can be studied in an organised way.
- Döbereiner grouped elements into triads — sets of three with similar properties in which the atomic weight of the middle element is the arithmetic mean of the other two.
- Newlands arranged the elements in increasing order of atomic weight and found that every eighth element, beginning from any element, resembles the first.
- Mendeleev’s periodic law states that the properties of elements are the periodic functions of their atomic weights. He placed 63 elements in vertical columns called groups and horizontal rows called periods, and left gaps for elements not yet discovered.
- Mendeleev’s table could not explain the positions of certain elements, the rare earth metals or the isotopes.
- Atomic number is the number of protons in the nucleus of an atom, and equals the number of electrons in a neutral atom. It is unique to an element, and it fixes the number and arrangement of electrons on which all properties depend.
- The modern periodic law, given by Moseley, states that the physical and chemical properties of elements are the periodic functions of their atomic number.
- Niels Bohr gave the long form of the modern periodic table. A periodic table is a tabular arrangement of elements in groups and periods, highlighting the regular trends in their properties.
- The table has eighteen vertical columns, the groups. A group is determined by the number of electrons in the outermost shell, so members of a group have similar properties.
- Group 1 are the alkali metals (except hydrogen), group 2 the alkaline earth metals, groups 3 to 12 the transition elements, group 16 the chalcogens, group 17 the halogens and group 18 the noble or inert gases.
- Groups 1, 2 and 13 to 18 are the main group, representative or normal elements. Their outermost shell is incomplete, except in group 18.
- The table has seven horizontal rows, the periods. A period is determined by the number of shells, so the period number equals the shell count. The periods hold 2, 8, 8, 18, 18, 32 and 32 elements.
- The lanthanides belong to group 3 of the sixth period and the actinides to group 3 of the seventh period. Both are shown below the table because they are large in number.
- The third period elements — Na, Mg, Al, Si, P, S, Cl and Ar — summarise the properties of their respective groups and are called typical elements.
- Periodic properties are those that reappear, or vary gradually, at regular intervals; the phenomenon is periodicity. Its cause is the recurrence of similar electronic configurations, and the recurrence appears after differences of 2, 8, 8, 18, 18 and 32 in atomic number.
- Down a group the number of shells increases one by one; across a period it stays the same.
- Valency is the combining capacity of an atom, equal to the number of electrons it can donate, accept or share. For 1, 2, 3 or 4 valence electrons the valency is 1, 2, 3 or 4; for 5, 6 or 7 it is \(8-5\), \(8-6\) and \(8-7\). Valency is always positive.
- Valency stays the same down a group. Across a period it rises to 4 at group 14 and then falls to 1 at group 17, and is 0 for group 18.
- The six periodic properties studied are atomic size, metallic character, non-metallic character, ionisation potential, electron affinity and electronegativity.
- Atomic size is the distance from the centre of the nucleus to the outermost shell, or half the inter-nuclear distance between two combined atoms. It is measured in angstrom, \(1\ \text{Å}=10^{-10}\ \text{m}\), or picometre, \(1\ \text{pm}=10^{-12}\ \text{m}\).
- Atomic size increases down a group, because shells are added, and decreases across a period, because the nuclear charge rises while the shell count stays the same.
- The inert gases are an exception: their atoms are bigger than the halogens of the same period, because the complete outermost shell has maximum electronic repulsion. Helium has the smallest size in the periodic table.
- A cation is smaller than its parent atom, since the remaining electrons are pulled in more strongly. An anion is larger, since the extra electrons repel one another and reduce the effective nuclear charge.
- Among isoelectronic ions, which have equal numbers of electrons, the greater the nuclear charge the smaller the ion.
- Metals tend to lose valence electrons and form positive ions. Metallic character increases down a group and decreases across a period. Metals are good reducing agents, and hydrogen, which also loses its single electron, is a good reducing agent too.
- Non-metals tend to gain electrons to attain the octet and form negative ions. Non-metallic character decreases down a group and increases across a period. Non-metals are oxidising agents.
- Across a period the oxides change from strongly basic through amphoteric to acidic. Down a group the basic nature of the oxides of metals increases.
- Across a period, chemical reactivity first decreases and then increases, being least in the middle. Down a group, the reactivity of metals increases and that of non-metals decreases.
- Melting and boiling points of metals decrease down a group, while those of non-metals increase. Across a period they usually increase upto group 14 and then decrease. Density increases gradually down a group, and across a period rises to a maximum before a slight decrease.
- Ionisation potential is the minimum energy required to remove the most loosely bound electron from a neutral isolated gaseous atom, converting it into a positive ion: \[\mathrm{M}\text{(g)} + \text{I.E.} \rightarrow \mathrm{M^{+}}\text{(g)} + e^{-}\] It is measured in eV/atom, with the S.I. unit \(\text{kJ mol}^{-1}\).
- Ionisation energy increases across a period, with exceptions, and decreases down a group. Helium has the highest value, \(2372.0\ \text{kJ mol}^{-1}\). Metals have low ionisation energy and non-metals high.
- Electron affinity is the energy released when a neutral isolated gaseous atom gains an electron to form an anion: \[\mathrm{X}\text{(g)} + e^{-} \rightarrow \mathrm{X^{-}}\text{(g)} + \text{E.A.}\] It is written with a negative sign; for chlorine it is \(-349\ \text{kJ mol}^{-1}\).
- Electron affinity increases across a period, being highest for group 17 and least for group 1, and decreases down a group. Fluorine and oxygen are exceptions, having lower values than chlorine and sulphur because their atoms are very small and the electron repulsions strong.
- Electronegativity is the tendency of an atom in a molecule to attract the shared pair of electrons towards itself. It is dimensionless. On the Pauling scale of 1932, fluorine is highest at 4.0 and caesium lowest at 0.7.
- Electronegativity increases across a period and decreases down a group. Metals are electropositive and non-metals electronegative, and oxidising power increases across a period and decreases down a group.
- The elements of the second period differ in properties from the other members of their groups, because of their small size and high electronegativity.
- Because group trends and period trends run in opposite directions, an element resembles the element diagonally below and to its right. This diagonal relationship links Li and Mg, Be and Al, and B and Si, which are called bridge elements.
- Mass number is the sum of the protons and the neutrons in the nucleus, so the neutron count is \[n = A - Z\] An \(n/p\) ratio near 1 means a stable nucleus, while a ratio of 1.5 and above means a radioactive one.
- The alkali metals have one valence electron, are soft reactive metals, good conductors, strong reducing agents, and react with water and acids to liberate hydrogen: \[2\,\mathrm{M}\text{(s)} + 2\,\mathrm{H_2O}\text{(l)} \rightarrow 2\,\mathrm{MOH}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
- The halogens have seven valence electrons, are coloured poisonous diatomic non-metals, non-conductors, and strong oxidising agents. They form electrovalent compounds with metals and covalent compounds with hydrogen and other non-metals.
Table 1.22 — Variation in the periodic properties across a period
| Property | Variation | Reason |
|---|---|---|
| Atomic size | Decreases → | Number of shells remains the same; nuclear charge increases |
| Metallic character | Decreases → | Atomic radius decreases; nuclear charge increases |
| Non-metallic character | Increases → | Atomic radius decreases; nuclear charge increases |
| Ionisation potential | Increases → | Atomic radius decreases; nuclear charge increases |
| Electron affinity | Increases → | Atomic radius decreases; nuclear charge increases |
| Electronegativity | Increases → | Atomic radius decreases; nuclear charge increases |
Table 1.23 — Variation in the periodic properties down a group
| Property | Variation | Reason |
|---|---|---|
| Atomic size (radii) | Increases ↓ | Number of shells increases |
| Metallic character | Increases ↓ | Atomic radii increase |
| Non-metallic character | Decreases ↓ | Atomic radii increase |
| Ionisation potential | Decreases ↓ | Atomic radii increase |
| Electron affinity | Decreases ↓ | Atomic radii increase |
| Electronegativity | Decreases ↓ | Atomic radii increase |
Atomic number increases down a group, but the effect of the nuclear charge is dominated by the increase in atomic radii, that is, by the increase in the number of shells.
Diagram
Fig. 1.9-1 — Trends In The Periodic Table At A Glance Outline grid of the periodic table with no symbols, drawn in the centre of the panel. Arrows placed around it: Across the top, two long arrows pointing right, labelled ‘ionisation potential increases’ and ‘electronegativity increases’. Across the grid diagonally, two banded arrows pointing right and up, labelled ‘non-metallic character increases’ and, pointing left and down, ‘metallic character increases’. Along the right edge, two vertical arrows pointing up, labelled ‘electronegativity increases’ and ‘ionisation potential increases’. Along the left edge, two vertical arrows pointing down, labelled ‘atomic radius increases’ and ‘electron affinity decreases’. Across the bottom, one long arrow pointing left labelled ‘atomic radius increases’ and one pointing right labelled ‘electron affinity increases’. Keep every arrow outside the grid so the grid stays clean, and use one colour for period trends and a second for group trends, with a small key.
1.10 Key Terms
- Actinides — the radioactive elements belonging to group 3 of the seventh period, shown separately below the main table.
- Alkali metals — the elements of group 1 (except hydrogen), which form strong alkalis with water.
- Alkaline earth metals — the elements of group 2, which form weaker alkalis than the group 1 elements.
- Amphoteric oxide — an oxide that behaves as a base towards an acid and as an acid towards a base.
- Atomic number — the number of protons in the nucleus of an atom, equal to the number of electrons in a neutral atom.
- Atomic size (atomic radius) — the distance between the centre of the nucleus of an atom and its outermost shell.
- Bridge elements — the second period elements Li, Be and B, each of which resembles the element diagonally below and to its right.
- Chalcogens — the elements of group 16, the oxygen family; the name means ore forming.
- Diagonal relationship — the resemblance between an element and the element placed at its right diagonal in the next period.
- Electron affinity — the energy released when a neutral isolated gaseous atom gains an electron to become an anion.
- Electronegativity — the tendency of an atom in a molecule to attract the shared pair of electrons towards itself.
- Electrovalent compound — a compound formed by the complete transfer of electrons from one atom to another.
- Group — a vertical column of the periodic table, determined by the number of electrons in the outermost shell.
- Halogens — the elements of group 17, whose name means salt formers.
- Ionisation potential — the minimum energy required to remove the most loosely bound electron from a neutral isolated gaseous atom.
- Isoelectronic ions — ions having the same number of electrons as one another.
- Lanthanides — the rare earth elements belonging to group 3 of the sixth period, shown separately below the main table.
- Law of octaves — Newlands’ observation that, in order of increasing atomic weight, every eighth element resembles the first.
- Main group elements — the elements of groups 1, 2 and 13 to 18, also called representative or normal elements.
- Mass number — the sum of the number of protons and the number of neutrons in the nucleus of an atom.
- Mendeleev’s periodic law — the statement that the properties of elements are the periodic functions of their atomic weights.
- Metallic character — the tendency of an element to lose its valence electrons and form a positive ion.
- Modern periodic law — the statement that the physical and chemical properties of elements are the periodic functions of their atomic number.
- Noble gases — the elements of group 18, whose outermost shell is complete and which hardly react.
- Non-metallic character — the tendency of an element to gain electrons in order to attain an octet in its outermost orbit.
- Nuclear charge — the positive charge in the nucleus of an atom, equal to the number of protons, that is, to the atomic number.
- Period — a horizontal row of the periodic table, determined by the number of shells in an atom.
- Periodic properties — properties that reappear, or vary gradually, at regular intervals in the periodic table.
- Periodic table — a tabular arrangement of elements in groups and periods that highlights the regular trends in their properties.
- Periodicity — the phenomenon of properties repeating at regular intervals, caused by the recurrence of similar electronic configurations.
- Radioactive element — an unstable element that emits radiations, having a neutron to proton ratio of about 1.5 or above.
- Reducing agent — a substance that loses electrons to another substance in a reaction.
- Shell (orbit) — one of the definite circular paths, counted outward from the nucleus, in which electrons revolve.
- Transition elements — the metallic elements of groups 3 to 12, which have their two outermost shells incomplete.
- Triad — a set of three elements with similar properties in which the atomic weight of the middle element is the mean of the other two.
- Typical elements — the third period elements from sodium to argon, which summarise the properties of their respective groups.
- Valence electrons — the electrons present in the outermost shell of an atom.
- Valence shell — the outermost shell of an atom, the shell whose electrons take part in reactions.
- Valency — the combining capacity of an atom, equal to the number of electrons it can donate, accept or share.
1.11 Points To Ponder
- Mendeleev’s table described the elements that were known and also predicted ones that were not. Does a classification become more trustworthy when it makes predictions that turn out right — and what would it have meant for his law if gallium had never been found?
- Ionisation energy rises across a period, yet beryllium and nitrogen sit above the line in period 2, and magnesium and phosphorus do the same in period 3. When an exception repeats itself in exactly the same place in every period, is it still an exception, or is it a sign that the rule is only part of the story?
- The elements at the middle of a period are called metalloids because they are neither clearly metals nor clearly non-metals. What does the existence of such a boundary suggest about treating ‘metal’ and ‘non-metal’ as two separate boxes rather than two ends of one scale?
- Helium is stable with two electrons in its outermost shell, while the other noble gases are stable with eight. Is the octet the real reason a noble gas is unreactive, or is it a convenient summary of something deeper about filled shells?
- Astatine and francium are so radioactive that their properties can barely be measured, yet the table still assigns them a place, a valency and a set of expected trends. How much confidence should be placed in properties that have been reasoned out from a position rather than measured in a laboratory?
1.12 Objective Questions
Multiple Choice Questions
Choose the correct option.
- The element having the smallest atomic size in the whole periodic table is: [1]
- hydrogen
- helium
- fluorine
- lithium
- An atom whose electronic configuration is 2, 8, 8, 2 belongs to: [1]
- period 3, group 2
- period 4, group 2
- period 2, group 4
- period 4, group 12
- The oxide of a third period element that behaves both as an acid and as a base is: [1]
- \(\mathrm{Na_2O}\)
- \(\mathrm{MgO}\)
- \(\mathrm{Al_2O_3}\)
- \(\mathrm{SO_3}\)
- On moving down group 17, the quantity that increases is the: [1]
- electronegativity
- ionisation potential
- atomic radius
- chemical reactivity
- Among \(\mathrm{Mg^{2+}}\), \(\mathrm{Na^{+}}\), \(\mathrm{F^{-}}\) and \(\mathrm{O^{2-}}\), the ion with the largest size is: [1]
- \(\mathrm{Mg^{2+}}\)
- \(\mathrm{Na^{+}}\)
- \(\mathrm{F^{-}}\)
- \(\mathrm{O^{2-}}\)
- The energy released when a neutral isolated gaseous atom accepts an electron is called: [1]
- ionisation potential
- electron affinity
- electronegativity
- nuclear charge
- The pair of elements showing a diagonal relationship is: [1]
- Li and Na
- Be and Al
- Na and Mg
- B and C
- The number of neutrons in an atom of \(^{40}_{18}\mathrm{Ar}\) is: [1]
- 18
- 20
- 22
- 40
- Elements placed in the same group of the periodic table have the same: [1]
- number of shells
- atomic size
- number of valence electrons
- mass number
- Of the elements listed below, the least reactive member of the third period is: [1]
- sodium
- silicon
- sulphur
- chlorine
Assertion–Reason
Each of the following gives an Assertion (A) and a Reason (R). Choose the answer from the codes below. (1) Both A and R are true and R is the correct explanation of A. (2) Both A and R are true but R is not the correct explanation of A. (3) A is true but R is false. (4) A is false but R is true.
Assertion (A): Atomic size decreases on moving from left to right in a period. Reason (R): The number of shells decreases from left to right in a period. [1]
Assertion (A): Alkali metals are strong reducing agents. Reason (R): They lose their single valence electron readily to attain a stable configuration. [1]
Assertion (A): An anion is larger than the atom from which it is formed. Reason (R): The added electron increases the repulsion between the electrons and reduces the effective nuclear charge. [1]
Assertion (A): The ionisation energy of caesium is lower than that of lithium. Reason (R): Ionisation energy decreases down a group because the atomic size increases. [1]
Match The Following
- Match each item in Column 1 with the correct item in Column 2. [5]
| Column 1 | Column 2 |
|---|---|
| a. 2, 8, 7 | i. chalcogens |
| b. group 16 | ii. lanthanides |
| c. helium | iii. highest ionisation energy |
| d. beryllium and aluminium | iv. bridge elements |
| e. group 3, period 6 | v. halogen of the third period |
Fill In The Blanks
- Complete the following statements. [4]
- The modern periodic law is based on the ________ of an element.
- Across a period from left to right, the metallic character ________.
- The S.I. unit of ionisation energy is ________.
- The valency of the elements of group 18 is ________.
True Or False
- State whether each of the following is true or false. [4]
- The transition elements are placed at the extreme right of the periodic table.
- The number of shells in an atom is equal to its period number.
- Electronegativity increases on moving down a group.
- Non-metals act as oxidising agents.
Name The Following
- Name each of the following. [4]
- The alkali metal present in period 4.
- The most electronegative element of the periodic table.
- The halogen that is a liquid at room temperature.
- The type of oxide formed by aluminium.
Give One Word Or Phrase For
- Give one word or phrase for each of the following. [4]
- The minimum energy needed to remove the most loosely bound electron from a neutral isolated gaseous atom.
- Ions that contain the same number of electrons.
- The elements of group 16 of the periodic table.
- Elements whose two outermost shells are incomplete.
Odd One Out, And Give A Reason
- Pick the odd one out in each set and give a reason for your choice. [3]
- Lithium, sodium, potassium, magnesium
- Fluorine, chlorine, bromine, argon
- Sodium, magnesium, aluminium, silicon
1.13 Check Your Concepts
Define valency. Why is valency always taken as a positive number? [2]
State the modern periodic law and name the scientist who put it forward. [2]
What is meant by periodicity of elements? State its cause. [2]
Give a reason why a cation is always smaller than the atom from which it is formed. [2]
Hydrogen is placed at the top of group 1, yet it is not counted among the alkali metals. Give two reasons. [2]
Explain why the atomic size decreases across a period even though the number of electrons in the atom is increasing. [3]
Describe how the nature of the oxides changes across the third period, naming one oxide of each kind. [3]
Differentiate between metallic character and non-metallic character, stating the basis of difference. [3]
Why does chemical reactivity first decrease and then increase on moving from sodium to chlorine in the third period? [3]
Differentiate between ionisation potential and electron affinity, stating the basis of difference, and state how each of them varies on moving down a group. [4]
What happens to the size of a chlorine atom when it gains one electron, and to the size of a magnesium atom when it loses two? Explain each change in terms of protons and electrons. [4]
Explain, in terms of atomic size and nuclear charge, why metallic character increases down group 1 but decreases across period 3. [4]
An element X has the electronic configuration 2, 8, 5. [5]
- State its group and period.
- State its valency.
- Is it a metal or a non-metal? Give a reason.
- Write the formula of its compound with hydrogen.
- Compare its atomic size with that of the element immediately to its right in the same period, and give a reason.
What is meant by a diagonal relationship? Name the three pairs of bridge elements, and explain why an element resembles the one placed diagonally below and to its right. [5]
1.14 Equations, Conversions & Identification
Write the balanced equation for the reaction of potassium with water. Name the gas liberated and state the nature of the solution formed. [3]
Complete and balance the following. [3]
- \(\mathrm{Na} + \mathrm{HCl} \rightarrow\)
- \(\mathrm{Mg} - 2e^{-} \rightarrow\)
- \(\mathrm{O} + 2e^{-} \rightarrow\)
Write the equations for the formation of a sodium ion from a sodium atom and of a chloride ion from a chlorine atom. State which of the two atoms acts as the reducing agent, and why. [3]
How would you obtain the following? In each case state the number of electrons lost or gained. [2]
- \(\mathrm{Mg^{2+}}\) from \(\mathrm{Mg}\)
- \(\mathrm{O^{2-}}\) from \(\mathrm{O}\)
Identify the element from the following observations: it lies in the third period, its atom has three valence electrons, and its oxide dissolves in both acids and alkalis. State its name, its symbol and its valency. [3]
With reference to Fig. 1.5-4, identify each of the following and state how its size compares with that of the parent atom. [2]
- The particle formed when a sodium atom loses one electron.
- The particle formed when a chlorine atom gains one electron.
With reference to Fig. 1.1-1, identify the following. [3]
- The group in which the salt formers are placed.
- The set of groups lying between group 2 and group 13, and the general name given to their elements.
- The two series of elements shown separately below the main body of the table.
Write the formula of the compound formed in each case, and state whether it is electrovalent or covalent. [3]
- Potassium and bromine
- Calcium and chlorine
- Hydrogen and iodine