Summary and Quick Revision

How to Use This Section

This is the last section of the chapter, meant to be read the night before the paper and again in the queue outside the hall.

Nothing new is taught here. Every card compresses something Sections 1 to 13 worked through properly, in the same notation and with the same numbers. If a line surprises you, go back and reread that section instead of memorising the line.

Eight cards, one mistake checklist, one 60-second panic list. Screenshot the three figures.

Key Point: The whole chapter is one idea. Properties of elements are periodic functions of their atomic numbers, because atomic number fixes the electronic configuration, and the outer configuration repeats after every 2, 8, 8, 18, 18, 32 elements. Every trend below — radius, ionization enthalpy, electron gain enthalpy, electronegativity, valence, oxide character — is the outer configuration plus two competing effects: nuclear charge pulling in and shielding by inner electrons pushing out. Across a period nuclear charge wins; down a group shielding and the new shell win. Hold on to that and the anomalies become footnotes instead of a separate list.

Card 1 — The Classification Timeline and Mendeleev's Predictions

Who did what, and where it broke

Key Point (Definition): Mendeleev's periodic law — the physical and chemical properties of the elements are a periodic function of their atomic weights. Modern periodic law — the physical and chemical properties of the elements are periodic functions of their atomic numbers. Write the correct noun; that is the one-mark difference.

Mendeleev's predictions

"Eka" is Sanskrit for one: eka-aluminium means "one place below aluminium" in the same group (group 13). Eka-boron was scandium, eka-manganese was technetium.

Element counts: about 31 known in 1800, 63 by 1865 (Mendeleev's table), 114 with official names when your textbook was printed and 118 today. Classification became necessary because nobody can remember 100-plus elements one at a time.

[NEET] Why Mendeleev is remembered and Newlands is not: Mendeleev trusted the pattern more than the data — he left gaps, predicted the properties of the missing elements, and even suggested that some atomic weights had been measured wrongly. Newlands forced every element into an octave whether it fit or not.

Card 2 — Structure of the Long Form of the Periodic Table

Periods — how long, and why

Key Point: The period number equals the principal quantum number $n$ of the valence shell, and the number of elements in a period is twice the number of orbitals being filled (one orbital, two electrons).

The third period has 8, not 18, because after $3p$ the next orbital in energy is $4s$, not $3d$ (the $n + l$ rule, $4s$: $4 + 0 = 4$ versus $3d$: $3 + 2 = 5$). So $3d$ is filled in the fourth period (Sc to Zn, the first transition series), $4d$ in the fifth (Y to Cd), $5d$ in the sixth (La, Hf to Hg) and $6d$ in the seventh.

Groups

Elements of one group have the same number of valence electrons in the same type of orbital, which is why they behave alike: the alkali metals Li, Na, K, Rb, Cs, Fr are all $ns^1$ with $n = 2, 3, 4, 5, 6, 7$.

Lanthanoids and actinoids

They are printed separately only to keep the table from becoming 32 columns wide. Elements after uranium (92) are the transuranium elements, all man-made.

The period, group and block finder

[Board] The 100th element (fermium) is in period 7. The fifth period has 18 elements because $n = 5$ allows $5s$ (1), $4d$ (5) and $5p$ (3) orbitals to fill before $6s$: 9 orbitals, 18 electrons.

Card 3 — IUPAC Nomenclature for Elements With Z > 100

Why it exists

Both American and Soviet groups claimed element 104 (rutherfordium versus kurchatovium). To avoid such fights, IUPAC gives every new element a temporary systematic name built from the digits of its atomic number until the discovery is confirmed and an official name is agreed.

The digit roots

Key Point: Roots are joined in order of the digits, the name ends in -ium, and the symbol is the first letter of each root (three letters). Two spelling rules: a final "i" of bi or tri is dropped before "ium" (so 102 is unnilbium, not unnilbiium), and a double "nn" from enn + nil is written with a single "n" (so 190 would be unennilium).

Elements 101 to 118 — compact

Practice the mechanism on unfamiliar numbers, because that is what is asked: $Z = 120$ is unbinilium (Ubn), $Z = 121$ is unbiunium (Ubu), $Z = 130$ is untrinilium (Utn), $Z = 150$ is unpentnilium (Upn).

[NEET] Three one-liner facts: element 106 is named after Seaborg (the only element named after a living person at the time); 117 tennessine and 118 oganesson are the newest halogen and noble gas; the temporary symbol always has three letters, the official one has one or two.

Card 4 — The Four Blocks; Metals, Non-metals and Metalloids

The four blocks in one table

The exceptions to remember

Metals, non-metals and metalloids

Metallic character increases down a group (Be to Ba, C to Pb) and decreases across a period (Na to Cl); non-metallic character does the opposite. The most metallic elements are at the bottom left (Cs, Fr), the most non-metallic at the top right (F, O).

[JEE/NEET] Because the change is gradual, Sn and Pb are metals, Ge and Si are metalloids, and C is a non-metal — all in the same group 14. Do not call boron a metal even though it heads group 13.

Card 5 — Atomic Radius and Ionic Radius

The definitions

Key Point (Definition): Atomic size cannot be measured directly (the electron cloud has no sharp edge), so it is defined from bond lengths. Covalent radius = half the distance between the nuclei of two bonded atoms of the same element in a molecule: the Cl-Cl distance in $\mathrm{Cl_2}$ is 198 pm, so $r(\mathrm{Cl}) = 99$ pm. Metallic radius = half the internuclear distance between neighbouring atoms in a metal crystal: the Cu-Cu distance is 256 pm, so $r(\mathrm{Cu}) = 128$ pm. Noble gases have only a van der Waals radius (from non-bonded contact), which is why their large values are not compared with the others in a period.

The data

The two rules and why

Ionic radius

Same element, different charge: the higher the positive charge, the smaller the ion — $\mathrm{Fe} > \mathrm{Fe^{2+}} > \mathrm{Fe^{3+}}$; $\mathrm{O^{2-}} > \mathrm{O^-} > \mathrm{O}$.

Isoelectronic species — the one rule

Key Point: Isoelectronic species have the same number of electrons. Among them, the one with the largest nuclear charge is the smallest, because the same cloud is pulled by more protons.

$$\mathrm{N^{3-}} > \mathrm{O^{2-}} > \mathrm{F^-} > \mathrm{Ne} > \mathrm{Na^+} > \mathrm{Mg^{2+}} > \mathrm{Al^{3+}} \qquad (\text{10 electrons; } Z = 7, 8, 9, 10, 11, 12, 13)$$

Other sets to recognise on sight: $\mathrm{S^{2-}} > \mathrm{Cl^-} > \mathrm{Ar} > \mathrm{K^+} > \mathrm{Ca^{2+}}$ (18 electrons); $\mathrm{H^-}$, He, $\mathrm{Li^+}$, $\mathrm{Be^{2+}}$ (2 electrons); $\mathrm{Br^-}$, Kr, $\mathrm{Rb^+}$, $\mathrm{Sr^{2+}}$ (36 electrons).

[NEET] Ordering traps: (1) $\mathrm{K^+}$ (18 electrons, $n = 3$ outer shell) is larger than $\mathrm{Na^+}$ (10 electrons, $n = 2$) — a new shell beats charge. (2) $\mathrm{Al^{3+}} < \mathrm{Mg^{2+}} < \mathrm{Na^+}$ — same shell, rising charge. (3) The largest atom in any period is the alkali metal, and the largest bonded atom overall is Cs (262 pm) or Fr — leave the noble gases out of size comparisons unless the question says van der Waals radius.

Card 6 — Ionization Enthalpy

The definition and its three conditions

Key Point (Definition): The first ionization enthalpy $\Delta_i H$ is the enthalpy change when the most loosely bound electron is removed from an isolated gaseous atom in its ground state: $$\mathrm{X(g)} \longrightarrow \mathrm{X^+(g)} + e^-, \qquad \Delta_i H_1$$ The second, $\Delta_i H_2$, removes the next electron from $\mathrm{X^+(g)}$. Units $\mathrm{kJ\ mol^{-1}}$. Every ionization enthalpy is positive (energy is always needed to pull an electron away from a positive nucleus), and $\Delta_i H_1 < \Delta_i H_2 < \Delta_i H_3$ always, because each electron leaves a more positive ion.

Three words that carry marks: isolated (no neighbouring atoms), gaseous (no lattice or solvent energy), ground state (not excited).

The periodicity graph, $Z = 1$ to 60

Maxima at the noble gases (He highest of all, 2372 $\mathrm{kJ\ mol^{-1}}$), minima at the alkali metals (Cs lowest, then Rb, K). Each period starts at a minimum (Li, Na, K, Rb, Cs) and climbs to a maximum (Ne, Ar, Kr, Xe), with two dips on the way in periods 2 and 3.

The factors on which it depends

The two classic anomalies

The order in period 2 is $\mathrm{Li} < \mathrm{B} < \mathrm{Be} < \mathrm{C} < \mathrm{O} < \mathrm{N} < \mathrm{F} < \mathrm{Ne}$, and in period 3 $\mathrm{Na} < \mathrm{Al} < \mathrm{Mg} < \mathrm{Si} < \mathrm{S} < \mathrm{P} < \mathrm{Cl} < \mathrm{Ar}$. Predicting aluminium from its neighbours — Na 496, Mg 737, Si 786 — gives "between 496 and 786 but below Mg", and the value is 577.

The group 13 irregularity

Ga is slightly above Al because the ten $3d$ electrons between Al and Ga shield poorly, so Ga's outer electron feels a larger effective nuclear charge than the size increase would suggest. Tl is above In for the same reason with the $4f$ electrons added (lanthanoid contraction).

Successive ionization enthalpies — the jump rule

Key Point: A sudden big jump between $\Delta_i H_n$ and $\Delta_i H_{n+1}$ means the $(n+1)$th electron comes from a noble-gas core. The element has $n$ valence electrons, sits in group $n$ (s-block) or $10 + n$ (p-block), and its stable valence is $n$.

The Na/Mg puzzle

First ionization: Na (loses its lone $3s$ electron) is lower than Mg (two $3s$ electrons, higher nuclear charge, smaller atom). Second ionization: Na is now $\mathrm{Na^+}$ with the stable neon core $2s^2 2p^6$, so the next electron is torn from a closed shell in a lower level — enormous cost. $\mathrm{Mg^+}$ still has one $3s$ electron to give. So $\Delta_i H_1(\mathrm{Na}) < \Delta_i H_1(\mathrm{Mg})$ but $\Delta_i H_2(\mathrm{Na}) \gg \Delta_i H_2(\mathrm{Mg})$.

[JEE Main] Hydrogen's ionization enthalpy in J per atom is $2.18 \times 10^{-18}$ J; multiply by Avogadro's number and you get $1312\ \mathrm{kJ\ mol^{-1}}$, or $1.312 \times 10^6\ \mathrm{J\ mol^{-1}}$ if the question asks in those units.

Card 7 — Electron Gain Enthalpy and Electronegativity

Electron gain enthalpy — definition and sign

Key Point (Definition): Electron gain enthalpy $\Delta_{eg} H$ is the enthalpy change when an electron is added to an isolated gaseous atom to form a gaseous anion: $$\mathrm{X(g)} + e^- \longrightarrow \mathrm{X^-(g)}, \qquad \Delta_{eg} H$$ It is negative (exothermic) for most elements because a neutral atom usually attracts an electron; it is positive (endothermic) for the noble gases (the electron must enter the next higher shell) and for a few $ns^2$ elements (Be, Mg) and $np^3$ elements (N) with stable configurations.

Sign convention: electron affinity (EA) is the energy released, so a positive EA means a negative $\Delta_{eg} H$ of roughly the same size: EA of Cl $= +349$ means $\Delta_{eg} H(\mathrm{Cl}) = -349\ \mathrm{kJ\ mol^{-1}}$. (Strictly $\Delta_{eg} H = -\mathrm{EA} - \tfrac{5}{2}RT$; ignore the small term unless asked.)

The data

Trends and the two anomalies

So chlorine has the most negative electron gain enthalpy of all elements ($-349$), not fluorine. Among the halogens the order is $\mathrm{Cl} > \mathrm{F} > \mathrm{Br} > \mathrm{I}$ (most to least negative). Fluorine is still the most electronegative element and the most reactive non-metal; those are different properties.

Electronegativity — the Pauling scale

Key Point (Definition): Electronegativity is the ability of an atom in a chemical compound to attract the shared electrons of a bond to itself. It is not a measurable quantity and has no units; it is a number on an arbitrary scale. On Pauling's scale fluorine is assigned 4.0 and every other element is placed relative to it. It depends on the bonding partner and the oxidation state, so it is a property of an atom in a molecule, not of an isolated atom.

Trend: increases across a period (Li 1.0 to F 4.0), decreases down a group (F 4.0 to At 2.2; Li 1.0 to Cs 0.7). It moves with ionization enthalpy and non-metallic character and against atomic radius and metallic character — smaller atom, higher $Z_{eff}$, stronger pull on bond electrons. The most electronegative element is F, the least is Cs (Fr).

Three properties that students mix up

[JEE Main] The Mulliken scale is $\chi_M = (IE + EA)/2$ in eV; divide by 2.8 to get a Pauling-like value. The second electron gain enthalpy of O and S ($\mathrm{O^-} \rightarrow \mathrm{O^{2-}}$) is strongly positive because a negative ion is being asked to accept another electron.

Card 8 — Chemical Trends: Valence, Hydrides and Oxides, Second-Period Anomalies, Reactivity and Oxide Character

Valence of the representative elements

Key Point: Valence $=$ number of valence (outermost) electrons, or $8 -$ that number. Groups 1, 2, 13, 14 use the first; groups 15, 16, 17 mostly the second (with respect to hydrogen).

Oxidation state follows electronegativity: in $\mathrm{OF_2}$ oxygen is $+2$ (F is more electronegative), in $\mathrm{Na_2O}$ oxygen is $-2$. Transition elements and actinoids show variable valence. In $[\mathrm{AlCl(H_2O)_5}]^{2+}$ the oxidation state of Al is $+3$ but its covalency is 6 (six bonds) — oxidation state and covalency are not the same number.

Anomalous behaviour of the second-period elements (Li, Be, B, C, N, O, F)

Diagonal relationship

Key Point: The first element of a group often resembles the second element of the next group because their sizes and charge/radius ratios are similar: Li and Mg, Be and Al, B and Si.

Chemical reactivity — the ends-of-period rule

Metallic character and reactivity of metals go down and left; non-metallic character and reactivity of non-metals go up and right. The elements change from electropositive (electron losers) at the left to electronegative (electron gainers) at the right.

Oxide character

Across period 3: $\mathrm{Na_2O}$ (strongly basic), MgO (basic), $\mathrm{Al_2O_3}$ (amphoteric), $\mathrm{SiO_2}$ (weakly acidic), $\mathrm{P_4O_{10}}$ (acidic), $\mathrm{SO_3}$ (strongly acidic), $\mathrm{Cl_2O_7}$ (most acidic). Down a group the oxides become more basic ($\mathrm{N_2O_3}$ acidic, $\mathrm{Bi_2O_3}$ basic). The higher the oxidation state of the same element, the more acidic the oxide ($\mathrm{SO_2} < \mathrm{SO_3}$, $\mathrm{Mn_2O_7}$ acidic while MnO is basic).

[Board] "Predict the formulas of the compounds between": silicon and bromine gives $\mathrm{SiBr_4}$ (valence 4 and 1); aluminium and sulphur gives $\mathrm{Al_2S_3}$ (3 and 2); element with $Z = 114$ and hydrogen gives $\mathrm{FlH_4}$ (group 14). Read the group, read the valence, cross-multiply.

The Mistakes That Cost the Most Marks

Each of these was flagged somewhere in Sections 1 to 13, ordered roughly by how often they show up in answer scripts.

1. Saying fluorine has the most negative electron gain enthalpy. It is chlorine ($-349\ \mathrm{kJ\ mol^{-1}}$); F is $-328$. The small $2p$ subshell of F repels the incoming electron. F is the most electronegative (4.0) and the most reactive non-metal — three different properties, three different winners.

2. Ranking B above Be and O above N in ionization enthalpy. Be ($2s^2$) > B ($2s^2 2p^1$): the $p$ electron is less penetrating and better shielded. N ($2p^3$, half-filled) > O ($2p^4$, one pair repelling). Period 2 order: $\mathrm{Li} < \mathrm{B} < \mathrm{Be} < \mathrm{C} < \mathrm{O} < \mathrm{N} < \mathrm{F} < \mathrm{Ne}$. Same shape in period 3 with Al/Mg and S/P.

3. Writing "atomic mass" in the modern periodic law, or "atomic number" in Mendeleev's. Mendeleev: periodic function of atomic weights. Modern (after Moseley, 1913): periodic function of atomic numbers. Moseley plotted $\sqrt{\nu}$ against $Z$ and got a straight line.

4. Calling helium s-block, or hydrogen an alkali metal. He is $1s^2$ but sits in group 18 because its shell is complete. H is placed separately — one electron like group 1, one short of a full shell like group 17.

5. Comparing a noble-gas radius with the others in its period. Noble-gas radii are van der Waals radii (non-bonded), so they come out larger than the halogen before them. Leave them out of the "smallest atom in the period" answer — in period 2 it is F (64 pm), not Ne.

6. Ordering an isoelectronic series by size the wrong way. Same electrons, so more protons means smaller: $\mathrm{O^{2-}} > \mathrm{F^-} > \mathrm{Na^+} > \mathrm{Mg^{2+}} > \mathrm{Al^{3+}}$. The anion with the lowest $Z$ is the largest. For the same element, $\mathrm{Fe} > \mathrm{Fe^{2+}} > \mathrm{Fe^{3+}}$.

7. Reading the p-block group number as the number of $p$ electrons. Group $= 10 +$ valence electrons for the p-block: $ns^2 np^4$ is group 16, not 4 or 6. For the d-block, group $=$ $(n-1)d$ electrons $+$ $ns$ electrons: $3d^6 4s^2$ is group 8. Period $=$ highest $n$.

8. Treating electronegativity as a measurable energy with units. Electronegativity is a relative, unitless number on an arbitrary scale (Pauling: F $= 4.0$) for an atom in a compound. Ionization enthalpy and electron gain enthalpy are measured for isolated gaseous atoms in $\mathrm{kJ\ mol^{-1}}$.

9. Getting the electron affinity sign backwards. Electron affinity is energy released, so a positive EA is a negative $\Delta_{eg} H$: EA(Cl) $= +349$, $\Delta_{eg} H(\mathrm{Cl}) = -349\ \mathrm{kJ\ mol^{-1}}$. Noble gases have positive $\Delta_{eg} H$ (He $+48$, Ne $+116$).

10. Confusing the second ionization enthalpy of Na and Mg. $\Delta_i H_1$: Na (496) < Mg (737). $\Delta_i H_2$: Na (about 4562) $\gg$ Mg (1451), because $\mathrm{Na^+}$ has the neon core. The big jump in successive values tells you the number of valence electrons and hence the group.

11. Missing the group 13 irregularity. Ionization enthalpy does not fall smoothly: B 801, Al 577, Ga 579, In 558, Tl 589. Poor shielding by $3d$ (for Ga) and $4f$ (for Tl) electrons keeps the outer electron more tightly held.

12. Calling $\mathrm{Al_2O_3}$ acidic or basic, and CO acidic. $\mathrm{Al_2O_3}$ (and $\mathrm{As_2O_3}$, ZnO, BeO) is amphoteric — reacts with both acids and bases. CO, NO and $\mathrm{N_2O}$ are neutral. Acidic: $\mathrm{Cl_2O_7}$, $\mathrm{SO_3}$, $\mathrm{CO_2}$; basic: $\mathrm{Na_2O}$, MgO. Across a period oxides go basic to acidic; down a group they become more basic.

Key Point: Two more that cost single marks: writing the IUPAC name with the roots in the wrong order or a four-letter symbol (105 is unnilpentium, Unp, official dubnium, Db), and forgetting that Zn, Cd and Hg are in the d-block but are not transition elements because their $d^{10}$ subshell is full in the element and its ions.

The 60-Second Revision

The irreducible minimum, for the queue outside the hall.

History. Dobereiner 1829 triads (Li-Na-K, Ca-Sr-Ba, Cl-Br-I; middle weight is the mean). de Chancourtois 1862 cylinder. Newlands 1865 octaves, works only to Ca. Lothar Meyer atomic-volume curves. Mendeleev 1869: periodic law by atomic weight, 63 elements, swapped Te/I, gaps for eka-Al (Ga: 68 predicted, 70 found; $\mathrm{E_2O_3}$, $\mathrm{ECl_3}$) and eka-Si (Ge: 72 predicted, 72.6 found; $\mathrm{EO_2}$, $\mathrm{ECl_4}$). Moseley 1913: $\sqrt{\nu} \propto Z$, modern periodic law by atomic number. Seaborg: actinoids, transuranium elements, seaborgium (106).

Table. 7 periods with 2, 8, 8, 18, 18, 32, 32 elements — twice the orbitals filled ($1s$; $2s2p$; $3s3p$; $4s3d4p$; $5s4d5p$; $6s4f5d6p$; $7s5f6d7p$). Period $=$ highest $n$. 18 groups: 1 alkali ($ns^1$), 2 alkaline earth ($ns^2$), 3 to 12 transition ($(n-1)d^{1-10}ns^{0-2}$), 13 to 18 p-block ($ns^2np^{1-6}$); 15 pnictogens, 16 chalcogens, 17 halogens, 18 noble gases. Lanthanoids Ce to Lu ($4f$), actinoids Th to Lr ($5f$), group 3. p-block group $= 10 +$ valence electrons; d-block group $= d + s$ electrons. He in group 18 (full shell); H alone. Zn, Cd, Hg: d-block, not transition. Metals 78%; metalloids B, Si, Ge, As, Sb, Te on the zig-zag.

IUPAC. nil un bi tri quad pent hex sept oct enn (0 to 9); roots in digit order + "ium"; three-letter symbol. 104 Unq/Rf, 106 Unh/Sg, 110 Uun/Ds, 114 Uuq/Fl, 117 Uus/Ts, 118 Uuo/Og, 120 Ubn.

Radius. Covalent (half of Cl-Cl 198 pm gives 99), metallic (half of Cu-Cu 256 gives 128), van der Waals for noble gases (do not compare). Decreases across (Li 152 to F 64; Na 186 to Cl 99), increases down (Li 152, Na 186, K 231, Rb 244, Cs 262; F 64, Cl 99, Br 114, I 133, At 140). Cation smaller (Na 186, $\mathrm{Na^+}$ 95), anion larger (F 64, $\mathrm{F^-}$ 136). Isoelectronic: more protons, smaller — $\mathrm{O^{2-}} > \mathrm{F^-} > \mathrm{Na^+} > \mathrm{Mg^{2+}} > \mathrm{Al^{3+}}$.

Ionization enthalpy. Isolated gaseous ground-state atom; always positive; $\Delta_i H_1 < \Delta_i H_2$. Increases across, decreases down; max at noble gases (He 2372), min at alkali metals. Anomalies Be > B, N > O (Mg > Al, P > S). Period 2: Li < B < Be < C < O < N < F < Ne. Na 496, Mg 737, Al 577, Si 786; B 801, Al 577, Ga 579, In 558, Tl 589. Big jump in successive values reveals the group. $\Delta_i H_2$: Na $\gg$ Mg. H: $1312\ \mathrm{kJ\ mol^{-1}}$.

Electron gain enthalpy. $\mathrm{X(g)} + e^- \to \mathrm{X^-(g)}$; more negative across, less negative down; noble gases positive (He +48, Ne +116, Ar +96). Cl $-349$ most negative of all; F $-328$; Br $-325$; I $-295$; O $-141$ < S $-200$. EA positive means $\Delta_{eg} H$ negative.

Electronegativity. Unitless, relative, atom in a bond, Pauling F = 4.0; Li 1.0, Be 1.5, B 2.0, C 2.5, N 3.0, O 3.5, F 4.0; Na 0.9 to Cl 3.0; Cs 0.7 lowest. Rises with $\Delta_i H$ and non-metallic character, falls with radius.

Chemistry. Valence = valence electrons or 8 minus them: hydrides LiH, $\mathrm{CaH_2}$, $\mathrm{B_2H_6}$, $\mathrm{CH_4}$, $\mathrm{NH_3}$, $\mathrm{H_2O}$, HF; oxides $\mathrm{Li_2O}$, MgO, $\mathrm{B_2O_3}$, $\mathrm{CO_2}$, $\mathrm{N_2O_5}$, $\mathrm{SO_3}$, $\mathrm{Cl_2O_7}$. Second period anomalous: tiny size, high charge/radius, high electronegativity, only 4 valence orbitals (covalency 4: $\mathrm{BF_4^-}$ versus $\mathrm{AlF_6^{3-}}$), $p\pi$-$p\pi$ bonds. Diagonal pairs Li/Mg, Be/Al, B/Si. Reactivity highest at both ends: metals down-left (Cs), non-metals up-right (F). Oxides: basic $\mathrm{Na_2O}$ (to NaOH), amphoteric $\mathrm{Al_2O_3}$, $\mathrm{As_2O_3}$, neutral CO, NO, $\mathrm{N_2O}$, acidic $\mathrm{Cl_2O_7}$ (to $\mathrm{HClO_4}$).

That is the whole chapter.