Classification of Elements and Periodicity — NEET notes

Historical Development of the Periodic Table

The periodic table didn't appear overnight — it evolved through Döbereiner's triads, Newlands' octaves, and Mendeleev's bold predictions. Mendeleev left gaps for undiscovered elements and was proved spectacularly right when Ga, Ge, and Sc were found. The modern periodic law (based on atomic number, not mass) fixed all remaining anomalies. NEET loves matching scientists to their contributions.

Modern Periodic Law: Properties = f(Z) , not f(A)

Döbereiner's Triads (1829)

Johann Döbereiner noticed that certain elements could be grouped in sets of three (triads) where the atomic mass of the middle element was approximately the arithmetic mean of the other two.

Döbereiner (1829) was one of the first to spot a pattern. He grouped elements into triads — sets of three where the middle element's atomic mass was roughly the average of the lightest and heaviest.

A_ middle A_ lightest + A_ heaviest 2

• Only 5 triads could be identified from the elements known at that time
• All three elements in a triad had similar chemical properties (e.g., Li, Na, K are all reactive metals)
• Not all elements could be grouped into triads — this was the main limitation
• This was the first serious attempt at classification, but it couldn't accommodate all known elements

Newlands' Law of Octaves (1866)

John Newlands arranged elements in order of increasing atomic mass and noticed that every 8th element had similar properties — analogous to the octave in music (do, re, mi, fa, sol, la, ti, do).

Newlands (1866) arranged all known elements by increasing atomic mass and noticed that the eighth element (starting from any given one) had properties similar to the first — like the eighth note in a musical scale repeating the first.

• Worked well only up to calcium (Z = 20)
• He assumed only 56 elements existed and no more would be discovered — wrong!
• He placed two elements in the same slot to force the pattern (e.g., Co and Ni together)
• No provision for noble gases (undiscovered at that time)
• Unlike Mendeleev, he did NOT leave gaps for undiscovered elements
• The law was ridiculed initially but recognized posthumously (Davy Medal, 1887)

Mendeleev's Periodic Table (1869)

Dmitri Mendeleev arranged 63 elements by increasing atomic mass and grouped them by similar chemical properties (especially oxide and hydride formulas). His genius was leaving gaps for undiscovered elements and predicting their properties.

Mendeleev (1869) stated: *'The properties of elements are a periodic function of their atomic masses.'* He arranged 63 known elements into a table with 8 groups and 12 periods, prioritizing chemical properties (oxide/hydride formulae) over strict mass order.

• Mendeleev's Periodic Law: Properties of elements are periodic functions of their atomic masses
• He changed the order of some elements to fit properties: Te (127.6) placed before I (126.9) — properties mattered more than mass
• He left gaps for eka-boron (Sc), eka-aluminium (Ga), and eka-silicon (Ge)
• Basis of classification: oxide and hydride formulae (e.g., R_2O_7 for halogens)
• Could accommodate noble gases later (as Group 0) without disrupting the table

Limitations of Mendeleev's Table

Despite its brilliance, Mendeleev's table had serious issues: isotopes couldn't be explained, anomalous pairs (Te/I, Co/Ni) broke mass ordering, and hydrogen's position was ambiguous.

Mendeleev's table was revolutionary, but several problems remained unresolved — and these 'cracks' ultimately pointed toward atomic number (not mass) as the true organizing principle.

• Position of hydrogen: Resembles both alkali metals (forms H^+ , like Na, K) and halogens (forms H^- , diatomic H_2 like Cl_2 ). No clear group assignment
• Anomalous pairs: Te (127.6) was placed before I (126.9) — violating increasing mass order. Same issue with Co (58.9) before Ni (58.7), Ar (39.9) before K (39.1)
• Isotopes: ^ 35 Cl and ^ 37 Cl have different masses — should they be placed separately? Mendeleev had no answer
• No place for lanthanides and actinides — they were awkwardly squeezed at the bottom
• Dissimilar elements in the same group: Cu and Ag were in the same group as alkali metals (sub-groups A and B)

Quick recall

Name three anomalous pairs in Mendeleev's table where a heavier element was placed before a lighter one.

Ar (39.9) before K (39.1), Co (58.9) before Ni (58.7), Te (127.6) before I (126.9). In all three cases, chemical properties demanded this order, but atomic mass didn't support it.

Modern Periodic Law (Moseley, 1913)

Henry Moseley showed that the fundamental property governing periodicity is the atomic number (Z), not atomic mass. This resolved all anomalies of Mendeleev's table.

Henry Moseley (1913) studied X-ray spectra of elements and discovered that the frequency of characteristic X-rays was related to atomic number (Z), not atomic mass. This led to the Modern Periodic Law.

= a(Z - b) (Moseley's equation)

• Moseley's contribution: linked X-ray frequency to atomic number experimentally
• Atomic number Z = number of protons = number of electrons (in neutral atom)
• Periodicity arises because electronic configurations repeat periodically with increasing Z
• The modern table has 7 periods and 18 groups (IUPAC numbering 1–18)

Modern Periodic Table — Structure & Electronic Configuration

The long form periodic table has 7 periods and 18 groups organized by electronic configuration. Elements fall into s, p, d, and f blocks based on which subshell is being filled. The period number equals the highest principal quantum number, and the group number relates directly to valence electron count. Understanding this structure is essential for predicting element properties in NEET.

Group number = cases n_s electrons & (s-block: Groups 1–2) 10 + n_s + n_p & (p-block: Groups 13–18) n_s + (n-1)d & (d-block: Groups 3–12) cases

Periods (Horizontal Rows)

There are 7 periods in the modern periodic table. The period number tells you the principal quantum number (n) of the outermost shell. Period length depends on how many electrons can fill orbitals at that energy level.

Each period corresponds to the filling of a new principal energy level. The period number equals the principal quantum number ( n ) of the outermost (valence) shell of the element.

• Period 1 is the shortest period (only 2 elements: H and He)
• Periods 2 and 3 are short periods (8 elements each)
• Periods 4 and 5 are long periods (18 elements each)
• Periods 6 and 7 are longest periods (32 elements each, including lanthanides/actinides)
• Maximum number of elements in a period = 2n^2 (but only for period 1; for higher periods, filling order matters)

Groups (Vertical Columns)

There are 18 groups in the IUPAC system. Elements in the same group have the same valence shell electronic configuration and hence similar chemical properties. Group number directly relates to valence electron count.

A group is a vertical column of elements with the same number and type of valence electrons. This is why elements in the same group show similar chemical behavior — their outermost electron arrangement is identical.

• IUPAC numbering: 1 to 18 (no IA/IIB confusion). NEET uses IUPAC numbering
• s-block: Groups 1 and 2 (+ He in some conventions)
• p-block: Groups 13 to 18
• d-block: Groups 3 to 12
• f-block: Lanthanides (4f) and Actinides (5f) — placed separately at the bottom
• Elements in the same group have the same general valence configuration → similar chemistry

s, p, d, f Blocks

The periodic table is divided into four blocks based on which subshell receives the differentiating (last) electron. This block classification determines the general properties of the element.

The block of an element is determined by the subshell into which the last (differentiating) electron enters. This simple rule divides the entire periodic table into four distinct regions.

• s-block: 2 groups × 7 periods = 14 elements. Includes H, alkali metals, alkaline earth metals, and He
• p-block: 6 groups × 6 periods = 36 elements (period 1 has no p-block). Contains all non-metals, metalloids, and some metals
• d-block: 10 groups × 4 series = 40 elements. All are metals with partially filled (n-1)d subshells
• f-block: 14 lanthanides ( 4f ) + 14 actinides ( 5f ) = 28 elements. Also called inner transition elements
• Helium ( 1s^2 ) is s-block by configuration but placed in Group 18 (p-block) due to its noble gas properties

Quick recall

An element has configuration [Xe]4f^ 14 5d^1 6s^2 . Which block does it belong to?

d-block. The differentiating (last) electron enters the 5d subshell. The 4f is completely filled. This is Lutetium (Lu). Despite having a filled 4f, it's classified as d-block because the last electron goes into 5d.

Predicting Position from Electronic Configuration

Given an element's electronic configuration, you can determine its period (= highest n), group, and block instantly. This is one of the most directly tested skills in NEET.

The periodic table is fundamentally a map of electronic configurations. Given the configuration, finding the position is mechanical:

• Period = highest principal quantum number ( n ) in the configuration
• Block = subshell of the last (differentiating) electron
• Group (s-block): = number of ns electrons (1 or 2)
• Group (p-block): = 10 + number of valence electrons ( ns + np )
• Group (d-block): = number of (n-1)d + ns electrons
• Group (f-block): No specific group number assigned (all are in Group 3)

Atomic & Ionic Radii

Atomic radius decreases across a period (increasing Zeff) and increases down a group (new shells). Cations are smaller than parent atoms, anions are larger. Isoelectronic species are compared by nuclear charge — more protons means smaller radius. These trends are among the most tested periodic properties in NEET.

Isoelectronic: r 1 Z_ eff More protons = Smaller radius

Types of Atomic Radii

Atoms don't have sharp boundaries, so 'radius' depends on how you measure it. The covalent radius is half the bond length in a homonuclear diatomic; van der Waals radius is half the distance between non-bonded atoms.

Since electron clouds don't have sharp edges, atomic 'size' is defined operationally based on the type of measurement.

• For any atom: r_ van der Waals > r_ metallic > r_ covalent
• Noble gases have only van der Waals radii (they don't form covalent bonds normally)
• The covalent radius is the most commonly used measure for comparison across the periodic table
• NCERT uses 'atomic radius' to generally mean covalent radius for bonded atoms

Trend Across a Period

As we move across a period, atomic number increases by 1 but electrons are added to the SAME shell. The effective nuclear charge (Zeff) increases, pulling the electron cloud inward, so atoms get smaller.

Moving left → right across a period: protons increase, electrons are added to the same shell (same n ). The increased nuclear charge pulls the electron cloud tighter → atomic radius decreases.

Z_ eff = Z - (Slater's rule: = shielding constant)

• Across a period: Z_ eff increases → atomic radius decreases
• Same principal quantum number ( n ) for all elements in a period → no new shell added
• Electrons in the same shell provide poor shielding for each other
• In Period 2: Li (152 pm) > Be (111 pm) > B (88 pm) > C (77 pm) > N (74 pm) > O (66 pm) > F (64 pm)
• Noble gas radii (van der Waals) are not directly comparable to covalent radii of other elements

Trend Down a Group

Going down a group, each element has one more electron shell than the one above. Despite increasing nuclear charge, the addition of a new shell dominates, and atomic radius increases.

Moving down a group: a new principal shell is added at each step. The outermost electrons are now farther from the nucleus, and inner shells shield them from nuclear attraction. The net effect: radius increases significantly.

• Down a group: new shell added → radius increases
• The increased nuclear charge is more than compensated by the shielding effect of inner electrons
• Group 1: Li (152 pm) < Na (186 pm) < K (227 pm) < Rb (248 pm) < Cs (265 pm)
• Group 17: F (64 pm) < Cl (99 pm) < Br (114 pm) < I (133 pm)
• The increase is large from Period 2→3 and 3→4, then more gradual due to d/f orbital shielding

Quick recall

Why does the atomic radius increase only slightly from Period 4 to Period 5 in transition metals?

Because of the poor shielding by 3d electrons. In Period 4, the 3d subshell fills before 4p. These 3d electrons don't shield the outer electrons effectively, so Z_ eff is higher than expected → radius doesn't increase as much from 4th to 5th period.

Ionic Radii

When an atom loses electrons (forms cation), it gets smaller because electron-electron repulsion decreases and Zeff per remaining electron increases. When it gains electrons (forms anion), the opposite happens.

Ions are NOT the same size as their parent atoms. Removing electrons (cation formation) shrinks the ion; adding electrons (anion formation) expands it.

• Cation < parent atom: Fewer electrons → less e⁻-e⁻ repulsion, same (or higher) nuclear charge per electron. Example: Na (186 pm) → Na⁺ (95 pm) — almost HALF!
• Anion > parent atom: More electrons → more e⁻-e⁻ repulsion, same nuclear charge spread over more electrons. Example: Cl (99 pm) → Cl⁻ (181 pm) — nearly DOUBLE!
• When an atom loses ALL valence electrons, the ion may lose an entire shell: Na ( 2,8,1 ) → Na⁺ ( 2,8 ) — the n=3 shell is gone
• Successive ionization: Na^+ > Mg^ 2+ > Al^ 3+ (all have 10 electrons, but nuclear charges increase: 11, 12, 13)

Isoelectronic Species

Isoelectronic species have the same number of electrons but different nuclear charges. Among them, the one with the most protons is the smallest because it exerts the greatest pull on the same number of electrons.

Isoelectronic species have the same total number of electrons. To compare their sizes, look at the nuclear charge (atomic number) — more protons means stronger pull on the same electron cloud → smaller radius.

For isoelectronic species: r 1 Z (more protons → smaller)

• Step 1: Count total electrons in each species
• Step 2: Confirm they're isoelectronic (same electron count)
• Step 3: Compare nuclear charges ( Z )
• Step 4: More protons → smaller radius. Arrange accordingly
• This method works for ANY set of isoelectronic species — not just the 10-electron series

Quick recall

Arrange in increasing order of ionic radius: K^+ , Ca^ 2+ , Cl^- , S^ 2-

All have 18 electrons. Nuclear charges: S(16), Cl(17), K(19), Ca(20). More protons → smaller. So: Ca^ 2+ < K^+ < Cl^- < S^ 2- (increasing radius).

Ionization Enthalpy

Ionization enthalpy (IE) is the energy required to remove the outermost electron from a gaseous atom. It increases across a period and decreases down a group. The famous anomalies — Be > B and N > O — arise from extra stability of fully-filled and half-filled subshells. Successive IEs increase sharply when a core shell is breached. NEET tests trends, anomalies, and numerical comparisons.

IE_1: M(g) M^+(g) + e^- ( H > 0, always endothermic)

Definition & Factors Affecting IE

Ionization enthalpy depends on nuclear charge, atomic radius, shielding effect, and the stability of the electronic configuration (half-filled/fully-filled extra stability).

Ionization enthalpy (IE₁) is the minimum energy needed to remove the most loosely bound electron from an isolated gaseous atom in its ground state. The atom must be in the gaseous state (no intermolecular forces) and ground state (no excited electrons).

M(g) _ i H_1 M^+(g) + e^-

• IE is always positive (endothermic) — energy must be supplied to remove an electron
• Units: kJ/mol (or eV per atom; 1 eV = 96.5 kJ/mol)
• Measured in gaseous state only — to eliminate intermolecular interactions
• Factors that INCREASE IE: Higher nuclear charge ( Z ), smaller atomic radius, less shielding, half-filled/fully-filled subshell stability
• Factors that DECREASE IE: Larger atomic radius, more shielding (inner shells), penetration of electron cloud

Trend Across a Period

IE increases across a period due to increasing Zeff. But there are two anomalies: Be > B (2s² > 2s²2p¹) and N > O (half-filled 2p³ > 2p⁴). These dips repeat in every period.

Across a period, Z_ eff increases steadily → IE should increase monotonically. But two anomalies break this trend in every period.

• General trend: IE increases across a period (Li < Be < B < C < N < O < F < Ne... with exceptions)
• Anomaly 1 — Be > B: Be ( 2s^2 , fully filled s) has higher IE than B ( 2s^2 2p^1 ). The lone 2p electron of B is easier to remove (higher energy, less penetration than 2s )
• Anomaly 2 — N > O: N ( 2p^3 , half-filled) has higher IE than O ( 2p^4 ). The 4th electron in O experiences electron-electron repulsion in a doubly occupied 2p orbital, making it easier to remove
• These anomalies repeat: Mg > Al (like Be > B) and P > S (like N > O)
• Noble gases have the highest IE in each period (fully filled configuration)

Trend Down a Group

Going down a group, atomic size increases and shielding increases. The outermost electron is farther from the nucleus and easier to remove, so IE decreases.

Down a group, each element has one more shell than the one above. The outermost electron is farther from the nucleus and screened by more inner electrons → easier to remove → IE decreases.

• Group 1 IE₁ (kJ/mol): Li (520) > Na (496) > K (419) > Rb (403) > Cs (376)
• Group 17 IE₁: F (1681) > Cl (1251) > Br (1140) > I (1008)
• The decrease is steepest from Period 2→3 (big jump in shell size)
• Alkali metals (Group 1) have the lowest IE₁ in each period
• Noble gases (Group 18) have the highest IE₁ in each period

Quick recall

Arrange in increasing order of IE₁: Na, Mg, K, Ca

Down the group: K < Na and Ca < Mg. Across the period: Na < Mg and K < Ca. Combined: K < Na < Ca < Mg (K is largest atom → lowest IE, Mg is smallest → highest IE).

Successive Ionization Enthalpies

IE₁ < IE₂ < IE₃ always. But there's a massive jump when you start removing electrons from a completely filled inner shell. This jump reveals the number of valence electrons.

Removing each subsequent electron costs more energy because you're pulling an electron from an increasingly positive ion. But the biggest clue is the sudden large jump in successive IEs.

IE_1 < IE_2 < IE_3 < (always increasing)

• Successive IEs always increase: each removal is from a more positive ion
• A disproportionately large jump signals that a new (inner) shell is being broken into
• The position of the big jump = number of valence electrons
• Na: jump after 1st → 1 valence electron. Mg: after 2nd → 2 VE. Al: after 3rd → 3 VE
• This is why Na forms Na⁺ (not Na²⁺) and Mg forms Mg²⁺ (not Mg³⁺) — the energy cost is too high beyond the valence shell

Electron Gain Enthalpy

Electron gain enthalpy (EGE) is the energy change when an electron is added to a gaseous atom. Halogens have the most negative values (love gaining electrons). Key anomalies: Cl is more negative than F (small size of F causes repulsion), and O is less negative than S. Noble gases have approximately zero (or positive) EGE. NEET tests sign conventions, trends, and anomalies.

M(g) + e^- M^-(g) _ eg H = negative (exothermic for most non-metals)

Definition & Sign Convention

Electron gain enthalpy is the enthalpy change when an extra electron is added to an isolated gaseous atom. If the process releases energy, EGE is negative; if energy must be supplied, it's positive.

Electron Gain Enthalpy (EGE) or _ eg H is the enthalpy change when an electron is added to an isolated gaseous atom in its ground state to form a gaseous anion.

M(g) + e^- M^-(g) _ eg H_1

• Negative _ eg H → energy is released → atom 'wants' the extra electron (exothermic)
• Positive _ eg H → energy must be supplied → atom resists gaining an electron (endothermic)
• Most non-metals have negative EGE (they form stable anions)
• Noble gases and elements with stable configurations (like Be: 2s^2 , N: 2p^3 , Mg: 3s^2 ) have positive or nearly zero EGE
• The more negative the EGE, the greater the atom's tendency to form an anion

Trends Across Period & Down Group

EGE becomes more negative across a period (atoms want electrons more as Zeff increases) and generally becomes less negative down a group (larger atoms hold extra electrons less tightly). But famous exceptions exist.

The general trend follows Z_ eff and atomic size, but the anomalies in EGE are more dramatic than those in IE.

• Across a period: EGE becomes more negative (→). Atoms become smaller, Z_ eff increases → incoming electron is held more tightly
• Down a group: EGE becomes less negative (↓). Atoms are larger → incoming electron is farther from nucleus, held less tightly
• Group 17 EGE (kJ/mol): F ( -328 ), Cl ( -349 ), Br ( -325 ), I ( -295 ). Cl is MORE negative than F!
• Group 16 EGE (kJ/mol): O ( -141 ), S ( -200 ), Se ( -195 ), Te ( -190 ). S is more negative than O!
• Noble gases have approximately zero (or slightly positive) EGE — their configurations are already fully stable

Special Cases

Noble gases, nitrogen, beryllium, and magnesium have near-zero or positive EGE because their electronic configurations are already highly stable (fully filled or half-filled subshells). Adding an electron would disrupt this stability.

Some elements actively resist gaining an extra electron. In these cases, EGE is zero or positive — you must supply energy to force the electron in.

• Noble gases: _ eg H > 0 (positive) — they are the most resistant to electron gain
• Be and Mg: _ eg H 0 or slightly positive — fully filled s-subshell
• N: _ eg H 0 — half-filled 2p³ stability (exchange energy is at maximum)
• These are the same elements that show anomalously high IE → same underlying cause (configuration stability)
• For 2nd electron gain enthalpy: O^-(g) + e^- O^ 2- (g) , _ eg H_2 > 0 (always positive, for ALL elements)

Electronegativity

Electronegativity measures an atom's ability to attract shared electrons in a chemical bond. Fluorine is the most electronegative element (4.0 on Pauling scale). It increases across a period and decreases down a group. The difference in electronegativity between bonded atoms determines bond polarity — ionic if difference > 1.7, covalent if < 0.4, polar covalent in between. NEET uses this to predict bond types and metallic character.

_F = 4.0 (Pauling scale) | > 1.7 ionic bond

Definition & Pauling Scale

Unlike IE and EGE (which are properties of isolated atoms), electronegativity describes an atom's electron-attracting power WITHIN a chemical bond. Pauling defined the scale using bond energy differences.

Electronegativity ( ) is the ability of an atom in a chemical bond to attract the shared pair of electrons toward itself. It's NOT a property of an isolated atom — it only has meaning when the atom is bonded.

• Pauling scale: Most widely used. Assigns F = 4.0 (highest), Cs = 0.7 (lowest among common elements)
• Mulliken scale: _M = IE + EA 2 (average of IE and electron affinity). More quantitative but less commonly used in NEET
• Electronegativity is dimensionless — it's a relative scale, not an absolute energy
• Noble gases are generally not assigned electronegativity values (they don't form bonds normally)
• It's a bond property — the same atom can have different effective EN in different bonding environments (though NEET treats it as fixed)

Trends Across Period & Down Group

Electronegativity increases across a period (smaller atoms with higher Zeff attract bonding electrons more) and decreases down a group (larger atoms hold bonding electrons less tightly).

The trends in electronegativity mirror those of IE and the inverse of atomic radius — smaller atoms with more nuclear charge attract bonding electrons more effectively.

• Across a period (L→R): EN increases. Example (Period 2): Li (1.0) < Be (1.5) < B (2.0) < C (2.5) < N (3.0) < O (3.5) < F (4.0)
• Down a group (T→B): EN decreases. Example (Group 17): F (4.0) > Cl (3.0) > Br (2.8) > I (2.5)
• Most electronegative: F (4.0) — top right of the periodic table (excluding noble gases)
• Least electronegative: Cs or Fr (~0.7) — bottom left of the periodic table
• The general rule: EN increases toward the top-right corner of the periodic table

Applications

The electronegativity difference between two atoms predicts bond type: non-polar covalent (Δχ ≈ 0), polar covalent (0.4 < Δχ < 1.7), or ionic (Δχ > 1.7). Also, low EN → metallic character, high EN → non-metallic character.

The difference in electronegativity ( ) between two bonded atoms is the simplest predictor of bond type. The larger the difference, the more ionic the bond.

• Metallic character ∝ 1 EN : Low EN → metals (easily lose electrons). High EN → non-metals (attract electrons)
• Metallic character increases down a group (EN decreases) and decreases across a period (EN increases)
• Most metallic element: Cs (or Fr). Most non-metallic element: F
• Diagonal relationship: elements diagonally placed (Li & Mg, Be & Al, B & Si) have similar EN → similar chemistry
• Pauling's formula: = 0.208 E where E is the extra bond energy in kJ/mol

Quick recall

Which bond is more polar: H–F or H–Cl?

H–F is more polar. (H–F) = 4.0 – 2.1 = 1.9. (H–Cl) = 3.0 – 2.1 = 0.9. Higher → more polar bond → H–F.

Chemical Reactivity Trends

Metallic character increases down a group and decreases across a period. Metal oxides are basic, non-metal oxides are acidic, and metalloid oxides are amphoteric. The diagonal relationship shows that Li~Mg, Be~Al, and B~Si have remarkably similar properties despite being in different groups. Reactivity of metals increases down Group 1 (easier to lose electrons) while reactivity of non-metals increases up Group 17 (higher EA). NEET tests oxide nature, reactivity patterns, and diagonal relationships.

Metallic character 1 IE 1 EN

Metallic vs Non-metallic Character

Metallic character depends on how easily an atom loses electrons (low IE → good metal). It increases down a group (IE decreases) and decreases across a period (IE increases). The zig-zag line in the periodic table separates metals, metalloids, and non-metals.

Metallic character = tendency to lose electrons and form cations. Non-metallic character = tendency to gain electrons and form anions. These are inversely related — as one increases, the other decreases.

• Across a period (L→R): Metallic character DECREASES (IE increases, EN increases). Na is a strong metal; Cl is a strong non-metal
• Down a group (T→B): Metallic character INCREASES (IE decreases, atoms can lose electrons more easily). Li is less metallic than Cs
• Metals: Low IE, low EN, lose electrons → form cations, conduct electricity, malleable, ductile
• Non-metals: High IE, high EN, gain electrons → form anions, insulators (except graphite), brittle
• Metalloids (B, Si, Ge, As, Sb, Te): Properties intermediate between metals and non-metals. They lie along the zig-zag line
• Most metallic: Cs (Fr). Most non-metallic: F

Nature of Oxides

The nature of element oxides changes systematically across a period: highly basic on the left (Na₂O), amphoteric in the middle (Al₂O₃), and acidic on the right (Cl₂O₇). This directly reflects the changing metallic/non-metallic character.

The oxide of an element reveals its metallic/non-metallic nature. Metal oxides are basic, non-metal oxides are acidic, and metalloid/borderline element oxides are amphoteric (react with both acids and bases).

• Across a period: oxide nature changes from basic → amphoteric → acidic
• Down a group: oxide nature becomes more basic (metallic character increases). Example: BeO (amphoteric) → BaO (strongly basic)
• Amphoteric oxides: Al_2O_3 , BeO , ZnO , PbO , SnO — react with both acids and bases
• Higher oxidation state → more acidic oxide. CrO (basic, Cr²⁺) vs CrO_3 (acidic, Cr⁶⁺)
• Acidic oxides are also called acid anhydrides (they form acids with water)

Reactivity of Metals & Non-metals

For metals, reactivity increases down a group (easier to lose electrons). For non-metals, reactivity increases up a group (higher tendency to gain electrons). This opposite trend is a beautiful consequence of the IE/EGE trends.

Here's one of the most elegant patterns in chemistry: metals and non-metals show opposite reactivity trends within their respective groups.

• Most reactive metal: Cs (or Fr) → lowest IE among stable elements
• Most reactive non-metal: F → highest EN, reacts with virtually everything
• Alkali metals: Li reacts slowly with water, Na reacts vigorously, K ignites, Cs explodes
• Halogens: F₂ is the strongest oxidizing agent, I₂ is relatively mild
• Gold (Au) and platinum (Pt) are among the least reactive metals → used in jewelry
• Noble gases are essentially unreactive → filled shells, no tendency to gain or lose electrons

Diagonal Relationship

Elements diagonally adjacent in the periodic table (e.g., Li & Mg, Be & Al) have similar properties because the effects of moving right (↑ EN, ↓ size) and moving down (↓ EN, ↑ size) roughly cancel out, giving similar charge-to-size ratios.

Moving one step right increases Z_ eff and decreases size. Moving one step down increases size and decreases Z_ eff . If you move diagonally (right AND down), these effects partially cancel, giving elements with similar charge densities and hence similar chemistry.

• Diagonal relationship is most prominent for Period 2 and Period 3 elements
• Cause: similar charge/radius ratio (polarizing power) of the diagonal pair
• Li ( = 1.0 ) ≈ Mg ( = 1.2 ) — similar EN. Be ( = 1.5 ) ≈ Al ( = 1.5 ) — identical EN!
• Li is the only alkali metal that forms a stable nitride (like Mg in Group 2)
• Be is the only alkaline earth metal with an amphoteric oxide (like Al in Group 13)
• This concept is unique to the top two periods — it fades for heavier elements

Quick recall

Why does Li resemble Mg more than Na (its own group member)?

Li and Mg have similar charge/radius ratios (polarizing power). Li⁺ is very small (high charge density) → more like Mg²⁺ than Na⁺. Both form nitrides, both carbonates decompose on heating (Na₂CO₃ is stable to heat), and both have lower solubility hydroxides.