3.3 Periodic Trends in Properties
How size, ionisation enthalpy, electron gain enthalpy and electronegativity change across periods and down groups, why beryllium, nitrogen and fluorine break the pattern, and how the same trends explain valence, the odd behaviour of second-period elements, diagonal relationships and the acid–base nature of oxides.
Sodium and chlorine are neighbours in the same period, yet sodium gives up an electron so easily that it is stored under oil, while chlorine snatches electrons from almost anything. Caesium, at the bottom of group 1, is even more reactive than sodium; fluorine, at the top of group 17, is more reactive than chlorine. These are not separate facts to memorise. They follow from a few properties of atoms — their size, how strongly they hold their outer electrons, and how strongly they attract extra ones — and those properties change in a regular, periodic way across the table.
This section studies the four key periodic properties: atomic and ionic radius, ionisation enthalpy, electron gain enthalpy and electronegativity. Each trend is explained by two competing effects — the attraction of the nucleus, measured by the effective nuclear charge, and the distance of the outer electrons, set by the number of shells — together with the extra stability of filled and half-filled subshells from Section 2.4. Then we use the trends to explain valence, the anomalous behaviour of the second-period elements, diagonal relationships, and why oxides on the left of the table are basic and those on the right acidic.
Atomic and Ionic Radii
An atom has no sharp edge — its electron cloud fades away gradually (Section 2.3) — so "the radius of an atom" must be defined from the distance between nuclei in a real substance, which X-ray diffraction measures very accurately.
- The covalent radius is half the distance between the nuclei of two identical atoms joined by a single covalent bond. In $\mathrm{Cl_{2}}$ the bond length is $198\ \mathrm{pm}$, so $r_{\mathrm{Cl}}=99\ \mathrm{pm}$.
- The metallic radius is half the distance between the nuclei of neighbouring atoms in a metallic crystal. In copper metal the distance is $256\ \mathrm{pm}$, so $r_{\mathrm{Cu}}=128\ \mathrm{pm}$.
- The van der Waals radius is half the distance between the nuclei of two identical non-bonded atoms in neighbouring molecules (or noble-gas atoms) in the solid. It is always larger than the covalent radius, because non-bonded atoms do not overlap.
Two effects decide the size of an atom. Going across a period, electrons are added to the same shell while the nuclear charge rises by one each step. Electrons in the same shell screen each other only weakly, so the effective nuclear charge $Z_{\mathrm{eff}}$ — the net positive charge an outer electron actually feels after the inner electrons have shielded part of $Z$ — increases, and the outer electrons are pulled closer. Going down a group, each step adds a whole new shell further from the nucleus; the extra inner electrons shield the extra protons, so $Z_{\mathrm{eff}}$ changes little and the size grows.
- Across a period (left to right) atomic radius generally decreases, because $Z_{\mathrm{eff}}$ increases while the valence shell stays the same. Period 2: $\mathrm{Li}\ 152\gt\mathrm{Be}\ 111\gt\mathrm{B}\ 88\gt\mathrm{C}\ 77\gt\mathrm{N}\ 74\gt\mathrm{O}\ 66\gt\mathrm{F}\ 64\ \mathrm{pm}$.
- Down a group atomic radius increases, because the principal quantum number of the valence shell increases. Group 1: $\mathrm{Li}\ 152\lt\mathrm{Na}\ 186\lt\mathrm{K}\ 231\lt\mathrm{Rb}\ 244\lt\mathrm{Cs}\ 262\ \mathrm{pm}$.
- Noble gases are not compared with these values: only their van der Waals radii can be measured, and those are much larger.
When an atom loses electrons it shrinks; when it gains electrons it swells. A cation is smaller than its parent atom — it often loses its whole outer shell, and the remaining electrons are held by the same nuclear charge. An anion is larger — the added electrons repel each other and the existing ones, while the nuclear charge is unchanged. So $r(\mathrm{Na^{+}})=95\ \mathrm{pm}$ against $r(\mathrm{Na})=186\ \mathrm{pm}$, but $r(\mathrm{Cl^{-}})=181\ \mathrm{pm}$ against $r(\mathrm{Cl})=99\ \mathrm{pm}$. (Ionic radii in this course are Pauling's values.)
Ionisation Enthalpy
How tightly does an atom hold its outermost electron? We measure it by the energy needed to pull that electron right away from an isolated atom in the gas phase.
Ionisation enthalpy depends on the attraction between the nucleus and the electron being removed: it is larger when $Z_{\mathrm{eff}}$ is larger, when the electron is closer to the nucleus, and when the electron belongs to a particularly stable arrangement. The inner electrons shield (screen) the outer ones from the nucleus; s electrons penetrate closer to the nucleus than p electrons of the same shell, and so both shield better and are held more tightly.
- Across a period $\Delta_{i}H_{1}$ generally increases (higher $Z_{\mathrm{eff}}$, smaller atom), from a minimum at the alkali metal to a maximum at the noble gas.
- Down a group $\Delta_{i}H_{1}$ decreases: the outer electron is further away and better shielded. Group 1: $\mathrm{Li}\ 520\gt\mathrm{Na}\ 496\gt\mathrm{K}\ 419\gt\mathrm{Rb}\ 403\gt\mathrm{Cs}\ 376\ \mathrm{kJ\,mol^{-1}}$.
- Exceptions in a period: $\mathrm{Be}\gt\mathrm{B}$ and $\mathrm{Mg}\gt\mathrm{Al}$, because the electron removed from B or Al is a $p$ electron, which is shielded by the $ns^{2}$ pair and held less tightly than an $s$ electron; $\mathrm{N}\gt\mathrm{O}$ and $\mathrm{P}\gt\mathrm{S}$, because N and P have the extra-stable half-filled $p^{3}$, while in O and S one $p$ orbital holds a pair, and the repulsion between the paired electrons makes one of them easier to remove.
| Period 2 | Li | Be | B | C | N | O | F | Ne |
|---|---|---|---|---|---|---|---|---|
| $\Delta_{i}H_{1}$ / $\mathrm{kJ\,mol^{-1}}$ | 520 | 900 | 801 | 1086 | 1402 | 1314 | 1681 | 2081 |
| Period 3 | Na | Mg | Al | Si | P | S | Cl | Ar |
| $\Delta_{i}H_{1}$ / $\mathrm{kJ\,mol^{-1}}$ | 496 | 738 | 578 | 787 | 1012 | 1000 | 1251 | 1521 |
Values are rounded from spectroscopic ionisation energies ($1\ \mathrm{eV}$ per atom $=96.485\ \mathrm{kJ\,mol^{-1}}$); tables in different books differ by a few $\mathrm{kJ\,mol^{-1}}$ (e.g. Be is often quoted as $899$).
Electron Gain Enthalpy and Electronegativity
Ionisation enthalpy measures how hard it is to take an electron away. The opposite question — how much an atom "wants" an extra electron — is answered by the electron gain enthalpy.
- Across a period $\Delta_{eg}H$ generally becomes more negative (higher $Z_{\mathrm{eff}}$, smaller atom), reaching the most negative values at the halogens (group 17), which need just one electron to complete the noble-gas configuration.
- Down a group it generally becomes less negative, as the added electron enters a shell further from the nucleus. Group 17: $\mathrm{Cl}\ (-349)\lt\mathrm{F}\ (-328)\lt\mathrm{Br}\ (-325)\lt\mathrm{I}\ (-295\ \mathrm{kJ\,mol^{-1}})$, so chlorine — not fluorine — has the most negative value.
- Exceptions: (i) F is less negative than Cl, and O ($-141$) less negative than S ($-200$): in the very small $2p$ subshell the incoming electron is strongly repelled by the electrons already there, which outweighs the higher attraction. (ii) Noble gases, and Be, Mg (filled $ns^{2}$) have positive values; N (half-filled $2p^{3}$) has a value close to zero.
Ionisation enthalpy and electron gain enthalpy describe isolated atoms. In a molecule, what matters is how strongly an atom pulls the shared electrons of a bond towards itself.
Electronegativity links directly to metallic character (Section 3.2): elements of low electronegativity readily give up electrons and are metallic; those of high electronegativity attract electrons and are non-metallic. The difference in electronegativity between two bonded atoms tells you how polar the bond is — a key idea of Chapter 4.
Valence, Anomalous Second-Period Elements and the Nature of Oxides
The valence of a representative element is the number of electrons it uses in bonding. It is usually equal to the number of valence electrons, or to eight minus that number, and so it repeats period after period — a property Mendeleev used to build his groups. The oxidation states of the elements (Chapter 7) follow the same pattern.
| Group | 1 | 2 | 13 | 14 | 15 | 16 | 17 | 18 |
|---|---|---|---|---|---|---|---|---|
| Valence electrons | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 |
| Valence | 1 | 2 | 3 | 4 | 3, 5 | 2, 6 | 1, 7 | 0, 8 |
| Hydride (period 3) | $\mathrm{NaH}$ | $\mathrm{MgH_{2}}$ | $\mathrm{AlH_{3}}$ | $\mathrm{SiH_{4}}$ | $\mathrm{PH_{3}}$ | $\mathrm{H_{2}S}$ | $\mathrm{HCl}$ | — |
| Highest oxide (period 3) | $\mathrm{Na_{2}O}$ | $\mathrm{MgO}$ | $\mathrm{Al_{2}O_{3}}$ | $\mathrm{SiO_{2}}$ | $\mathrm{P_{4}O_{10}}$ | $\mathrm{SO_{3}}$ | $\mathrm{Cl_{2}O_{7}}$ | — |
The first element of each group, in period 2 (Li, Be, B, C, N, O, F), differs noticeably from the rest of its group. The reasons are its small size, large charge/radius ratio, high electronegativity and the absence of d orbitals in its valence shell. As a result, period-2 elements have a maximum covalence of $4$ (boron forms $\mathrm{[BF_{4}]^{-}}$, while aluminium forms $\mathrm{[AlF_{6}]^{3-}}$), and they readily form $p\pi$–$p\pi$ multiple bonds ($\mathrm{C{=}C}$, $\mathrm{C{\equiv}N}$, $\mathrm{N{\equiv}N}$, $\mathrm{O{=}O}$) that their heavier group members seldom form. Lithium and beryllium form compounds with marked covalent character, unlike the rest of groups 1 and 2.
Periodic trends also explain chemical reactivity. On the far left, elements with very low ionisation enthalpy (the alkali metals) lose electrons easily and are strong reducing agents; on the far right, elements with very negative electron gain enthalpy and high electronegativity (the halogens) gain electrons easily and are strong oxidising agents. Reactivity is therefore highest at the two ends of a period and lowest in the middle; among metals it increases down a group, among non-metals it decreases.
Across a period the oxides change from strongly basic (left) through amphoteric to strongly acidic (right); down a group the basic character increases. Period 3:
$\mathrm{Na_{2}O}$, $\mathrm{MgO}$ (basic) $\ \rightarrow\ $ $\mathrm{Al_{2}O_{3}}$ (amphoteric) $\ \rightarrow\ $ $\mathrm{SiO_{2}}$, $\mathrm{P_{4}O_{10}}$, $\mathrm{SO_{3}}$, $\mathrm{Cl_{2}O_{7}}$ (acidic)
A basic oxide reacts with water to form a base or with acids to form salts; an acidic oxide gives an acid with water or reacts with bases; an amphoteric oxide reacts with both acids and bases. A few non-metal oxides, such as $\mathrm{CO}$, $\mathrm{NO}$ and $\mathrm{N_{2}O}$, are neutral.
The s-block elements and the general trends of the p-block are taught and tested in school, but not in the year-end board exam. JEE Advanced lists both the s-block and the p-block. JEE Main and NEET list only the p-block — the general trends of groups 13–18 and the unique behaviour of the first element of each group; neither lists the s-block.
The s-block: groups 1 and 2. The alkali metals (Li, Na, K, Rb, Cs, Fr) have the outer configuration $ns^{1}$; the alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra) have $ns^{2}$. Their trends are the cleanest in the table (Interactive 3 below):
- Size increases down each group, as a new shell is added each period. A group-2 atom is smaller than its group-1 neighbour, because of its extra nuclear charge.
- Ionisation enthalpy falls down each group, so metallic character and reactivity increase: caesium reacts far more violently with water than lithium. Group 1 shows only the oxidation state $+1$; group 2 has higher ionisation enthalpies and shows only $+2$.
- The hydration enthalpy of the ions falls as they get larger: $\mathrm{Li^{+}}\gt\mathrm{Na^{+}}\gt\mathrm{K^{+}}\gt\mathrm{Rb^{+}}\gt\mathrm{Cs^{+}}$. The very strongly hydrated $\mathrm{Li^{+}}$ ion makes lithium the strongest reducing agent of the group in water, although its ionisation enthalpy is the highest.
- The oxides and hydroxides are basic, more strongly so down each group (Rule 3.10). Many of the metals and their salts colour a flame: Li crimson, Na yellow, K lilac, Ca brick red, Sr crimson, Ba apple green (Be and Mg give no colour).
Anomalous lithium and beryllium. The first member of each group is exceptionally small, with a high charge/radius ratio, so its compounds have marked covalent character. Lithium is harder and has a higher melting point than the other alkali metals; it burns in air to give mainly the normal oxide $\mathrm{Li_{2}O}$ (sodium gives the peroxide $\mathrm{Na_{2}O_{2}}$, potassium the superoxide $\mathrm{KO_{2}}$); it alone combines directly with nitrogen, forming $\mathrm{Li_{3}N}$; and $\mathrm{Li_{2}CO_{3}}$ decomposes on heating, whereas the other alkali metal carbonates do not. In these ways lithium resembles magnesium (Definition 3.10). Beryllium has a very small atom and high ionisation enthalpies, so its compounds are largely covalent and easily hydrolysed; with only four orbitals in its valence shell it never exceeds a coordination number of four (magnesium and the others can reach six); and its oxide and hydroxide are amphoteric, like those of aluminium.
The p-block: general trends in groups 13–18. The outer configuration is $ns^{2}np^{1}$ to $ns^{2}np^{6}$ (helium $1s^{2}$).
- Oxidation states. The highest (group) oxidation state equals the number of valence electrons: $+3$ in group 13 up to $+7$ in group 17. Non-metals also show negative states (N $-3$, O $-2$, Cl $-1$).
- Inert pair effect. In groups 13–15 a state two below the group state becomes more stable down the group, because the $ns^{2}$ pair of the heavy elements is reluctant to take part in bonding: Tl(I) is more stable than Tl(III), Pb(II) than Pb(IV) — so $\mathrm{PbO_{2}}$ is a strong oxidising agent — and Bi(III) than Bi(V).
- Metallic character increases down each group: carbon and silicon are non-metals, germanium a metalloid, tin and lead metals. Non-metals sit at the top right of the block.
- The first member is different (as in this section): small size, high electronegativity and no d orbitals give a maximum covalence of $4$ and strong $p\pi$–$p\pi$ multiple bonds — $\mathrm{N{\equiv}N}$ and $\mathrm{O{=}O}$ molecules, while phosphorus and sulphur form singly bonded $\mathrm{P_{4}}$ and $\mathrm{S_{8}}$.
- Irregularities come from poor shielding by filled d and f subshells: gallium (135 pm) is slightly smaller than aluminium (143 pm), and thallium's ionisation enthalpy is higher than indium's. The noble gases (group 18) have the highest ionisation enthalpies of their periods, falling down the group — which is why xenon, the easiest to ionise of the stable ones, forms compounds such as $\mathrm{XeF_{2}}$ and $\mathrm{XeF_{4}}$.
Worked example 1. (a) Arrange $\mathrm{Li^{+}}$, $\mathrm{Na^{+}}$ and $\mathrm{K^{+}}$ in order of decreasing hydration enthalpy. (b) Lithium has the highest first ionisation enthalpy in group 1 ($520\ \mathrm{kJ\,mol^{-1}}$, against $419$ for potassium), yet it is the strongest reducing agent in water. Explain.
Solution. (a) $\mathrm{Li^{+}}\gt\mathrm{Na^{+}}\gt\mathrm{K^{+}}$: the smallest ion has the highest charge density and attracts water molecules most strongly. (b) Turning the solid metal into hydrated ions involves atomisation, ionisation and hydration. Lithium needs more energy to ionise, but the huge hydration enthalpy of the tiny $\mathrm{Li^{+}}$ ion more than repays it, so overall lithium gives up its electron in water most readily.
Worked example 2. (a) Carbon dioxide is a gas of separate $\mathrm{O{=}C{=}O}$ molecules, but silica, $\mathrm{SiO_{2}}$, is a hard solid. (b) $\mathrm{PbO_{2}}$ oxidises concentrated hydrochloric acid to chlorine, but $\mathrm{CO_{2}}$ and $\mathrm{SiO_{2}}$ do not. Explain both using p-block trends.
Solution. (a) The small carbon atom forms strong $p\pi$–$p\pi$ double bonds to oxygen, so each C is satisfied by two O atoms in a small molecule. The larger Si atom forms poor π bonds and instead makes four single Si–O bonds, linking every atom into a giant network. (b) The inert pair effect: for lead, the $+2$ state is more stable than $+4$, so Pb(IV) readily gains two electrons (reduced to $\mathrm{Pb^{2+}}$), oxidising $\mathrm{Cl^{-}}$ to $\mathrm{Cl_{2}}$. For carbon and silicon the group state $+4$ is the stable one.
Exercises
- 1 Distinguish between covalent radius and van der Waals radius. Which is larger for chlorine?
- 2 The $\mathrm{Br{-}Br}$ bond length is $228\ \mathrm{pm}$. Find the covalent radius of bromine.
- 3 Explain why atomic radius decreases from Li to F.
- 4 Arrange K, Mg, Na, Rb in order of increasing atomic radius.
- 5 Why is a cation smaller and an anion larger than the parent atom?
- 6 Arrange $\mathrm{S^{2-}}$, $\mathrm{Cl^{-}}$, $\mathrm{K^{+}}$, $\mathrm{Ca^{2+}}$ in order of decreasing radius.
- 7 Which of the following are isoelectronic with $\mathrm{Ne}$: $\mathrm{Na^{+}}$, $\mathrm{Cl^{-}}$, $\mathrm{O^{2-}}$, $\mathrm{Mg^{2+}}$, $\mathrm{K^{+}}$, $\mathrm{F^{-}}$?
- 8 By what percentage is the radius of $\mathrm{Mg^{2+}}$ ($65\ \mathrm{pm}$) smaller than that of the Mg atom ($160\ \mathrm{pm}$)?
- 1 Define first ionisation enthalpy and write the equation for the first ionisation of magnesium.
- 2 Why are successive ionisation enthalpies of an element always larger than the previous one?
- 3 Arrange B, Be, C, O, N in order of increasing first ionisation enthalpy.
- 4 Explain why $\Delta_{i}H_{1}$ of phosphorus is greater than that of sulphur.
- 5 Why does $\Delta_{i}H_{1}$ decrease from Li to Cs?
- 6 An element has $\Delta_{i}H_{1}=738$, $\Delta_{i}H_{2}=1451$ and $\Delta_{i}H_{3}=7733\ \mathrm{kJ\,mol^{-1}}$. To which group does it belong? Explain.
- 7 Calculate the energy needed to convert $1.00\ \mathrm{g}$ of gaseous magnesium atoms into $\mathrm{Mg^{2+}}$ ions ($M=24.31\ \mathrm{g\,mol^{-1}}$).
- 8 Find the energy per atom (in J) needed to ionise a potassium atom ($\Delta_{i}H_{1}=419\ \mathrm{kJ\,mol^{-1}}$) and the longest wavelength of light that can do it.
- 1 Define electron gain enthalpy. What does a negative value mean?
- 2 Which has the most negative electron gain enthalpy of all the elements? Why not fluorine?
- 3 Why is the second electron gain enthalpy of oxygen positive?
- 4 Arrange O, S, F, Cl in order of increasingly negative electron gain enthalpy.
- 5 Why do Be and Mg have positive electron gain enthalpies, while Li and Na have negative ones?
- 6 How does electronegativity differ from electron gain enthalpy?
- 7 Arrange Na, Cl, Mg, P, Si in order of increasing electronegativity.
- 8 Find the electronegativity difference in $\mathrm{NaCl}$ and in $\mathrm{C{-}H}$ (Na $0.9$, Cl $3.0$, C $2.5$, H $2.1$). Which bond is more polar?
- 1 Predict the formulae of the compounds of (a) calcium and chlorine, (b) aluminium and sulphur, (c) silicon and oxygen.
- 2 Name the three diagonal pairs and give one property shared by Be and Al.
- 3 Give two reasons for the anomalous behaviour of the second-period elements.
- 4 Classify as acidic, basic, amphoteric or neutral: $\mathrm{MgO}$, $\mathrm{SO_{3}}$, $\mathrm{Al_{2}O_{3}}$, $\mathrm{CO}$, $\mathrm{K_{2}O}$, $\mathrm{P_{4}O_{10}}$.
- 5 Write a balanced equation for the reaction of $\mathrm{SO_{3}}$ with water and of $\mathrm{CaO}$ with water.
- 6 Why is reactivity highest at the two ends of a period?
- 7 Which is the stronger reducing agent, sodium or potassium? Which is the stronger oxidising agent, fluorine or iodine?
- 8 Applied. A water-treatment engineer must choose between $\mathrm{CaO}$ and $\mathrm{SO_{2}}$ to raise the pH of acidic mine drainage. Which, and why?
- 1 Arrange Li, Na, K, Rb and Cs in order of (a) increasing atomic radius, (b) decreasing first ionisation enthalpy, (c) decreasing hydration enthalpy of the $\mathrm{M^{+}}$ ion.
- 2 Give two ways in which lithium differs from the other alkali metals, and one way in which it resembles magnesium.
- 3 Explain why beryllium chloride is largely covalent and why beryllium never shows a coordination number above four.
- 4 Use the inert pair effect to explain why $\mathrm{PbCl_{2}}$ is more stable than $\mathrm{PbCl_{4}}$, and why thallium commonly forms $\mathrm{Tl^{+}}$ while aluminium forms only $\mathrm{Al^{3+}}$.
- 5 Nitrogen exists as $\mathrm{N_{2}}$ molecules, but phosphorus as $\mathrm{P_{4}}$ units with single bonds. Explain.
E1. NEET The correct order of radii is (A) $\mathrm{N^{3-}\lt O^{2-}\lt F^{-}\lt Na^{+}}$ (B) $\mathrm{Na^{+}\lt F^{-}\lt O^{2-}\lt N^{3-}}$ (C) $\mathrm{F^{-}\lt Na^{+}\lt O^{2-}\lt N^{3-}}$ (D) $\mathrm{Na^{+}\lt N^{3-}\lt O^{2-}\lt F^{-}}$
E2. JEE Main The first ionisation enthalpies of Na, Mg and Si are $496$, $738$ and $787\ \mathrm{kJ\,mol^{-1}}$. The value for Al is closest to (A) $575$ (B) $760$ (C) $800$ (D) $1000\ \mathrm{kJ\,mol^{-1}}$
E3. Board · 3 marks Explain why (i) the first ionisation enthalpy of nitrogen is higher than that of oxygen; (ii) chlorine has a more negative electron gain enthalpy than fluorine; (iii) $\mathrm{Na^{+}}$ is smaller than $\mathrm{Na}$.
E4. JEE Main Numerical answer. An element of period 3 has successive ionisation enthalpies $787$, $1577$, $3232$, $4356$ and $16\,091\ \mathrm{kJ\,mol^{-1}}$. What is its group number?
E5. JEE Advanced One or more options correct. (A) The first ionisation enthalpy of Be is greater than that of B. (B) Electronegativity decreases from F to I. (C) $\mathrm{Al_{2}O_{3}}$ reacts with both $\mathrm{HCl}$ and $\mathrm{NaOH}$. (D) Of all the elements, fluorine has the most negative electron gain enthalpy.
E6. Board · Case study The oxides of the period-3 elements are $\mathrm{Na_{2}O}$, $\mathrm{MgO}$, $\mathrm{Al_{2}O_{3}}$, $\mathrm{SiO_{2}}$, $\mathrm{P_{4}O_{10}}$, $\mathrm{SO_{3}}$ and $\mathrm{Cl_{2}O_{7}}$. (i) How does the acid–base nature of these oxides change across the period, and why? (2 marks) (ii) Name the amphoteric oxide. (1 mark) (iii) Write the equation for the reaction of $\mathrm{Na_{2}O}$ with water. (1 mark)
Research question. Why is lithium — and increasingly sodium — the metal of choice for rechargeable batteries in phones, electric cars and grid storage?
Why it matters. Lithium-ion batteries power billions of devices and most electric vehicles. Lithium is mined mainly in Australia, Chile, Argentina and China, and demand is rising so fast that companies in China, Europe and India are developing sodium-ion batteries, using a cheaper and far more abundant element from the same group.
What to do.
- Use the group-1 data in this section (atomic radius, $\Delta_{i}H_{1}$, electronegativity) to explain why the alkali metals give up an electron so readily.
- A battery stores more energy per kilogram if its ions are light. Compare the molar masses of Li ($6.94$), Na ($22.99$) and K ($39.10\ \mathrm{g\,mol^{-1}}$): how many moles of each metal are there in $1.00\ \mathrm{kg}$?
- In a battery the ions move in and out of solid electrodes. Using the ionic radii $\mathrm{Li^{+}}\ 60$, $\mathrm{Na^{+}}\ 95$, $\mathrm{K^{+}}\ 133\ \mathrm{pm}$ (Pauling), explain why a small ion is an advantage, and estimate by how much the volume of $\mathrm{Na^{+}}$ exceeds that of $\mathrm{Li^{+}}$.
- Find out (from an energy agency report or a recent review) one advantage and one drawback of sodium-ion batteries compared with lithium-ion.
Tools: Rule 3.7 and Rule 3.8, Interactive 2 (group 1), and a reliable energy or materials-science source.
Hand in: A short table of the group-1 data and your calculations, and a half-page recommendation on where sodium-ion batteries make most sense.
Going further. Lithium has the most negative standard electrode potential of all the elements, although its ionisation enthalpy is the highest in group 1. Find out what this has to do with the very large hydration enthalpy of the small $\mathrm{Li^{+}}$ ion (Chapter 7 and Class 12).
🧪 Self-check quiz
Q1. Which species is the smallest?
Q2. The first ionisation enthalpy of oxygen is lower than that of nitrogen because
Q3. Which element has the most negative electron gain enthalpy?
Q4. Going down group 17, electronegativity
Q5. Which oxide is amphoteric?