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Andhra Pradesh (BIEAP)Class 9 Physical Science★ AP Board (BSEAP) Class 9 Physical Science · ~8 marks · Board: AP Board

Chemical Bonding

Why atoms BOND — the OCTET RULE explained. IONIC BOND — NaCl and MgO formation, crystal lattice, properties. COVALENT BOND — H₂, O₂, N₂, CH₄, H₂O, CO₂. COORDINATE BOND. Lewis dot structures. Comparing ionic vs covalent. AP Board Class 9 Physical Science.

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1Explain why atoms form bonds using the octet/duet rule
2Describe ionic bond formation for NaCl, MgO, and CaCl₂ with electron dot diagrams
3List the properties of ionic compounds (MP/BP, conductivity, solubility, brittleness)
4Describe covalent bond formation with single, double, and triple bonds; give examples
5Draw Lewis dot structures for H₂, Cl₂, O₂, N₂, CH₄, H₂O, CO₂
6Compare ionic and covalent compounds in a table
7Explain coordinate (dative) bonds with NH₄⁺ as an example

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Why this chapter matters

Chemical Bonding directly follows from electronic configurations and the octet rule — the conceptual foundation is built in the Atomic Structure chapter. Ionic bond formation (showing electron transfer for NaCl, MgO, CaCl₂ with electron dot diagrams) is a standard 4-5 mark question. Covalent bond Lewis structures (H₂, O₂, N₂, CH₄) are tested in 4-5 mark questions. The ionic vs covalent properties comparison table is a reliable 3-4 mark question. This chapter feeds directly into Class 10 chapters (acids/bases, carbon compounds, metallurgy).

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AP Board Class 9 Physical Science: Atomic Structure (electronic configurations, octet rule)

Chemical Bonding — Class 9 Physical Science

"Atoms bond to become STABLE. The noble gases already are. Every other atom is trying to become one. This is the OCTET RULE — the central organizing principle of chemistry."

1. Why Do Atoms Bond?

Atoms bond to achieve a STABLE ELECTRONIC CONFIGURATION — usually 8 electrons in the outermost shell (OCTET). Exceptions: Hydrogen and Helium stable with 2 electrons (DUET). 'The noble gases (Group 18: He, Ne, Ar) have complete outer shells. They are UNREACTIVE. Every other atom, in some sense, is TRYING to be a noble gas.'

Three Ways to Achieve Stability

LOSE electrons (metals — become cations). 2. GAIN electrons (non-metals — become anions). 3. SHARE electrons (non-metals with non-metals).

Kossel-Lewis Approach

G.N. Lewis and W. Kossel independently proposed: atoms achieve stability by acquiring 8 electrons in the valence shell. Lewis introduced the OCTET RULE and DOT STRUCTURES.

2. Ionic (Electrovalent) Bond — Electron TRANSFER

Occurs between a METAL (low ionisation energy — readily loses electrons) and a NON-METAL (high electron affinity — readily gains electrons). Metal LOSES electron(s) → CATION (+). Non-metal GAINS electron(s) → ANION (−). Electrostatic ATTRACTION between opposite charges forms the BOND.

Formation of NaCl (Sodium Chloride)

Na (Z=11): 2, 8, 1 → loses 1e⁻ → Na⁺ (2, 8) — NEON configuration (stable). Cl (Z=17): 2, 8, 7 → gains 1e⁻ → Cl⁻ (2, 8, 8) — ARGON configuration (stable). Na⁺ + Cl⁻ → NaCl (held by ionic/electrostatic bond).

Formation of MgO (Magnesium Oxide)

Mg (Z=12): 2, 8, 2 → loses 2e⁻ → Mg²⁺ (2, 8). O (Z=8): 2, 6 → gains 2e⁻ → O²⁻ (2, 8). Mg²⁺ + O²⁻ → MgO.

Formation of CaCl₂ (Calcium Chloride)

Ca (Z=20): 2, 8, 8, 2 → loses 2e⁻ → Ca²⁺ (2, 8, 8). Each Cl (Z=17): 2, 8, 7 → gains 1e⁻ → Cl⁻. Ca²⁺ + 2Cl⁻ → CaCl₂. 'One Ca²⁺ ion attracts TWO Cl⁻ ions — the formula is CaCl₂, not CaCl.'

Properties of Ionic Compounds

HIGH melting and boiling points — strong electrostatic forces hold the crystal lattice. Much energy needed to break.
Conduct electricity when MOLTEN (liquid) or DISSOLVED IN WATER — ions FREE to move.
Do NOT conduct in SOLID state — ions LOCKED in fixed positions.
BRITTLE — when force shifts layers, like charges repel → crystal SHATTERS.
Usually SOLUBLE IN WATER (water molecules hydrate and separate ions).

Occurs between TWO NON-METALS with similar electronegativities. Neither transfers — they SHARE valence electrons. Each atom COUNTS the shared electrons toward its octet.

Types of Covalent Bonds

Bond Type	Shared Pairs	Example	Lewis Structure
Single	1 pair (2e⁻)	H₂, Cl₂, CH₄, H₂O	H−H, Cl−Cl
Double	2 pairs (4e⁻)	O₂, CO₂	O=O, O=C=O
Triple	3 pairs (6e⁻)	N₂	N≡N

Key Examples

H₂: Each H has 1 e⁻. Share 1 pair → both get 2 (duet). Cl₂: Each Cl has 7 valence e⁻. Share 1 pair → both get 8 (octet). O₂: Each O has 6 valence e⁻. Share 2 pairs → both get 8. N₂: Each N has 5 valence e⁻. Share 3 pairs → both get 8. 'N₂ has the STRONGEST bond — a triple bond with bond energy 941 kJ/mol.' CH₄ (Methane) : C has 4 valence e⁻, shares with 4 H atoms → each H gets 2, C gets 8. H₂O: O has 6 valence e⁻, shares with 2 H atoms → each H gets 2, O gets 8. CO₂: C has 4 valence e⁻, shares 2 pairs with each O → all atoms get 8.

Properties of Covalent Compounds

LOW melting and boiling points — weak INTERMOLECULAR forces between separate molecules. WITHIN-molecule bonds are strong. BETWEEN-molecule forces are weak.
Do NOT conduct electricity — no free ions or electrons (except graphite — delocalised electrons between layers).
Usually INSOLUBLE IN WATER — soluble in ORGANIC SOLVENTS.

4. Coordinate (Dative) Bond

A covalent bond where BOTH shared electrons come from the SAME atom. Once formed, it is identical to a regular covalent bond. Example — NH₄⁺ (Ammonium ion) : NH₃ has a LONE PAIR on nitrogen. H⁺ (no electrons) accepts this lone pair → N→H coordinate bond. Example — H₃O⁺ (Hydronium ion) : H₂O donates a lone pair to H⁺.

5. Ionic vs Covalent — Quick Comparison

Property	Ionic	Covalent
Formation	Electron TRANSFER	Electron SHARING
Between	Metal + Non-metal	Non-metal + Non-metal
MP/BP	HIGH	LOW
Conductivity	Only when molten/dissolved	NO (except graphite)
Solubility	Water	Organic solvents
Physical state	Crystalline solid	Gas, liquid, or soft solid

6. Common Mistakes to Avoid

'NaCl molecules exist independently' — NaCl is an IONIC CRYSTAL LATTICE. The formula represents the simplest RATIO (1:1) — not a discrete molecule.
'Covalent bonds are weak' — WITHIN-molecule covalent bonds are VERY STRONG. The INTERMOLECULAR forces are WEAK — that's why BP/MP are low.
'Ionic compounds never conduct electricity' — They DO conduct — but ONLY when MOLTEN or DISSOLVED (ions FREE to move). NEVER as solids.
'CaCl has the same formula as NaCl' — Ca²⁺ needs TWO Cl⁻ ions to balance charge. Formula is CaCl₂. Always cross-multiply valencies.

7. AP SSC Exam Focus

Topic	Marks	Type
Ionic bond formation (NaCl, MgO, CaCl₂)	4-5	Equation + explanation
Covalent bond examples (H₂, O₂, CH₄)	4-5	Lewis structures
Ionic vs Covalent properties	3-4	Table comparison
Coordinate bond (NH₄⁺)	2-3	MCQ or Short

Key Memory Aids

IONIC = Transfer (think: I Transferred money). Metal to Non-metal.
COVALENT = Sharing (think: Co-workers Share). Non-metal with Non-metal.
Ionic compounds: 'HIGH and MIGHTY' — HIGH MP/BP. MIGHT conduct when mobile (molten/dissolved).
Covalent compounds: 'LOW and LONELY' — LOW MP/BP. LONELY molecules (weakly attracted to each other).

Valency vs Ionic Charge

Valency is the COMBINING CAPACITY (no sign). Ionic charge has a SIGN. Na has valency 1, forms Na⁺ (charge +1). Cl has valency 1, forms Cl⁻ (charge −1). Mg has valency 2, forms Mg²⁺. O has valency 2, forms O²⁻. Al has valency 3, forms Al³⁺. N has valency 3, forms N³⁻. 'Use VALENCY when writing formulas (cross-multiply the numbers). Use IONIC CHARGE when explaining the bond (opposite charges attract).' Example: Al³⁺ and O²⁻ → cross-multiply valencies → Al₂O₃ (Al: 2 atoms, O: 3 atoms). Check charges: 2(+3) + 3(−2) = +6 −6 = 0 (correct).

Key formulas & results

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Ionic and Covalent Bonding

WHY ATOMS BOND: To achieve STABLE electronic configuration — usually OCTET (8 electrons in outermost shell). H and He stable with 2 (DUET). Noble gases (He, Ne, Ar) already stable — unreactive. THREE WAYS: LOSE electrons → cation (metals). GAIN electrons → anion (non-metals). SHARE electrons → covalent (non-metal + non-metal). IONIC BOND (ELECTROVALENT): Metal LOSES e⁻ → cation (+). Non-metal GAINS e⁻ → anion (−). Electrostatic attraction. EXAMPLES: NaCl: Na (2,8,1) − 1e⁻ → Na⁺ (2,8) [neon config]. Cl (2,8,7) + 1e⁻ → Cl⁻ (2,8,8) [argon config]. MgO: Mg (2,8,2) − 2e⁻ → Mg²⁺ (2,8). O (2,6) + 2e⁻ → O²⁻ (2,8). CaCl₂: Ca (2,8,8,2) − 2e⁻ → Ca²⁺. 2 × Cl + 1e⁻ each → 2 Cl⁻. VALENCY CROSS-MULTIPLY: Ca(2+) and Cl(1−) → Ca₁Cl₂ → CaCl₂. Al(3+) and O(2−) → Al₂O₃. IONIC PROPERTIES: HIGH MP/BP (strong electrostatic forces). Conducts ONLY when molten or dissolved (ions free). NO conduction in solid (ions locked). BRITTLE. Usually SOLUBLE in water. COVALENT BOND: Non-metal + non-metal, SHARE electrons. Single bond = 1 shared pair. Double = 2 pairs. Triple = 3 pairs. EXAMPLES: H₂ (H−H, single), Cl₂ (single), O₂ (O=O, double), N₂ (N≡N, triple — STRONGEST bond 941 kJ/mol), CH₄ (C with 4 single bonds), H₂O (O with 2 single bonds + 2 lone pairs), CO₂ (O=C=O, two double bonds). COVALENT PROPERTIES: LOW MP/BP (weak intermolecular forces). NO conductivity (no ions/free electrons — except graphite). Usually INSOLUBLE in water. Soluble in organic solvents. COORDINATE (DATIVE) BOND: Both shared electrons come from ONE atom. NH₃ donates lone pair to H⁺ → NH₄⁺. H₂O donates lone pair to H⁺ → H₃O⁺.

AP EXAM KEY TRAPS: (1) CaCl has wrong formula — should be CaCl₂ (Ca valency=2, Cl valency=1, cross-multiply). ALWAYS cross-multiply valencies to get formula. (2) Ionic compounds do NOT conduct in SOLID state (ions locked in lattice) — only when MOLTEN or DISSOLVED. (3) Covalent WITHIN-molecule bonds are STRONG. Between-molecule (intermolecular) forces are WEAK — that's why MP/BP are low. (4) NaCl exists as a crystal LATTICE, not discrete molecules. (5) BILE MEMORY AID: Ionic = TRANSFER (I Transferred), Covalent = SHARING (Co-workers Share).

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Common mistakes & fixes

These are the exact errors that cost students marks in board exams. Read them once, save yourself the trouble.

WATCH OUT

✗ Writing CaCl instead of CaCl₂, or AlO instead of Al₂O₃

✓ When writing ionic compound formulas, always use the CROSS-MULTIPLY rule with valencies (combining capacities): Ca has valency 2, Cl has valency 1. Cross-multiply: Ca₁Cl₂ → CaCl₂. Al has valency 3, O has valency 2. Cross-multiply: Al₂O₃. CHECK: The total positive charge must equal total negative charge. CaCl₂: Ca²⁺ provides +2, two Cl⁻ provide −2 total. Sum = 0 (correct). Al₂O₃: 2 × Al³⁺ = +6. 3 × O²⁻ = −6. Sum = 0 (correct). For NaCl: Na⁺ (+1) + Cl⁻ (−1) = 0 (correct). This charge neutrality rule is the fastest way to verify ionic formulas in the exam.

Practice problems

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5-minute revision

The whole chapter, distilled. Read this the night before the exam.

•WHY ATOMS FORM BONDS: To achieve STABLE noble gas configuration (8 electrons in outermost shell — OCTET RULE; 2 electrons for K shell — DUET RULE for H and He). Noble gases already have this stable configuration → unreactive.
•THREE BONDING MECHANISMS: (1) LOSE electrons → form positive ion (cation) — metals do this. (2) GAIN electrons → form negative ion (anion) — non-metals do this. (3) SHARE electrons → covalent bond — non-metal + non-metal.
•IONIC BOND (Electrovalent): Forms between METAL and NON-METAL. Metal LOSES electrons → cation (+). Non-metal GAINS electrons → anion (−). Electrostatic attraction between opposite charges holds them together. Example: Na (2,8,1) − 1e⁻ → Na⁺ (2,8 — neon configuration). Cl (2,8,7) + 1e⁻ → Cl⁻ (2,8,8 — argon configuration). The ions form a crystalline LATTICE (not discrete NaCl molecules).
•CROSS-MULTIPLY VALENCIES FOR FORMULA: Ca²⁺ + Cl⁻ → cross-multiply gives Ca₁Cl₂ → CaCl₂. Al³⁺ + O²⁻ → Al₂O₃. The rule ensures total positive charge = total negative charge (charge neutrality). NaCl: +1 + (−1) = 0 (correct). MgO: +2 + (−2) = 0 (correct).
•IONIC COMPOUND PROPERTIES: (1) HIGH melting/boiling points (strong electrostatic forces throughout lattice). (2) Conduct electricity ONLY when molten or dissolved (ions free to move). NO conduction in solid (ions locked in lattice). (3) Generally SOLUBLE in water (polar solvent). (4) Insoluble in organic solvents. (5) BRITTLE — when stressed, like-charges align and repel, shattering the crystal.
•COVALENT BOND: Forms between TWO NON-METALS. SHARE electron pairs. Single bond = 1 shared pair (H₂, Cl₂, CH₄). Double bond = 2 shared pairs (O₂, CO₂). Triple bond = 3 shared pairs (N₂ — strongest at 941 kJ/mol). Number of bonds = number of electrons each atom needs to complete its octet.
•COVALENT EXAMPLES — LEWIS STRUCTURES: H₂ (H−H). Cl₂ (Cl−Cl, each Cl has 3 lone pairs). O₂ (O=O, each O has 2 lone pairs). N₂ (N≡N, each N has 1 lone pair). CH₄ (C with 4 H, all single bonds — tetrahedral). H₂O (O with 2 H + 2 lone pairs — bent). CO₂ (O=C=O — linear).
•COVALENT COMPOUND PROPERTIES: (1) LOW melting/boiling points (weak inter-molecular forces, even though bonds within molecule are strong). (2) Do NOT conduct electricity (no ions or free electrons — exception: graphite). (3) Generally INSOLUBLE in water. (4) SOLUBLE in organic solvents. (5) Soft or gaseous at room temperature.
•COORDINATE (DATIVE) BOND: A special covalent bond where BOTH shared electrons come from ONE atom (the donor). The other atom (acceptor) contributes no electrons. Example: in NH₄⁺, three N-H bonds are normal covalent. The 4th bond between N and H⁺ is coordinate — both electrons come from N's lone pair. Similarly H₃O⁺ (H₂O + H⁺).
•IONIC vs COVALENT TABLE: 5 KEY PROPERTIES — (1) Bonding: transfer vs sharing. (2) Between which elements: metal+non-metal vs non-metal+non-metal. (3) MP/BP: high vs low. (4) Conductivity: in molten/aqueous vs none. (5) Solubility: water vs organic solvents.

Andhra Pradesh (BIEAP) marks blueprint

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Where this shows up in the real world

This chapter isn't just an exam topic — it lives in the world around you.

Salt production in AP coastal areas

Salt (NaCl) is produced industrially by evaporating sea water along AP's 974-km coastline. The salt pans of Krishnapatnam, Nellore, and other AP coastal areas produce thousands of tonnes annually. The ionic nature of NaCl explains both its solubility (allows extraction from sea water) and its high melting point (allows it to be solid at room temperature for transport). Understanding ionic bonding is the foundation of the salt industry.

Petrochemical industry and covalent bonding

Visakhapatnam's HPCL refinery and AP's petrochemical industry process hydrocarbons — carbon compounds held together by COVALENT bonds. Crude oil is a mixture of covalent hydrocarbon molecules. Distillation separates them by boiling point (which depends on molecular size and intermolecular forces — properties from this chapter). LPG, petrol, diesel, kerosene, and bitumen all have different covalent molecular structures. The Class 9 covalent bond concept underlies this entire industry.

Battery technology and ion movement

Every battery (alkaline, lithium-ion, lead-acid) works by ion movement — the principle that ionic compounds conduct only when ions can move. In a lithium-ion battery, Li⁺ ions move from one electrode to another through an electrolyte (which allows ion movement). Solid-state batteries (an emerging technology AP's research labs are studying) replace liquid electrolytes with solid ones that still permit ion movement. The Class 9 principle of ionic conduction is the foundation of all battery technology.

Exam strategy

Battle-tested tips from teachers and toppers for this chapter.

Ionic bond diagram (4-5 marks): for each compound (NaCl, MgO, CaCl₂), draw — (1) the electron dot structure of each atom showing valence electrons, (2) the transfer arrow(s), (3) the resulting cation and anion with charges. Show the configurations before and after. Each component earns marks.

Lewis structure (3-4 marks): for each covalent molecule, draw — (1) the atom symbols, (2) all bonding electron pairs as dots or lines, (3) all lone pairs explicitly. CH₄: C in centre, 4 H around, 4 single bonds. Don't forget lone pairs on N₂ (1 on each N), O₂ (2 on each O), and H₂O (2 on O).

Ionic vs covalent table (3-4 marks): always use a TABLE format — 5 rows × 2 columns. Property in left, ionic in middle, covalent in right. Five properties: how bond forms, between which elements, MP/BP, conductivity, solubility. Tables are clearer and faster than paragraphs.

Formula writing (1 mark each): for any ionic formula, use the cross-multiply rule. Write the valencies, swap as subscripts, verify charge neutrality. CaCl₂, Al₂O₃, Na₂O, MgCl₂.

Coordinate bond (2 marks): name the bond, explain that BOTH electrons come from one atom (the donor), give the NH₄⁺ example with NH₃ + H⁺ → NH₄⁺ showing the donation. Name the donor (N) and acceptor (H⁺).

Going beyond the textbook

For olympiad aspirants and curious learners — topics that build on this chapter.

STRETCH

Research electronegativity (Linus Pauling scale) — the tendency of an atom to attract a shared electron pair. The DIFFERENCE in electronegativity between two atoms determines bond type: large difference (>1.7) → ionic; medium difference → polar covalent; small/zero difference → nonpolar covalent. Pauling's electronegativity scale (1932) is one of the most useful concepts in chemistry. Research how it predicts molecular polarity and reactivity.

STRETCH

Investigate the structure of water and hydrogen bonding — water (H₂O) is a covalent molecule but its molecules attract each other through HYDROGEN BONDS — a special intermolecular force. This is why water has unusually high boiling point (100°C — way higher than predicted from molecular weight). Hydrogen bonding also explains why ice floats (less dense than liquid water), DNA structure (base pairing), and protein folding. Research hydrogen bonds in biology.

STRETCH

Explore band theory of solids — extending covalent bonding to crystals. In metals, valence electrons are DELOCALISED across the entire crystal — forming an 'electron sea' that explains metallic conductivity and ductility. In semiconductors (silicon, used in chips), the band structure allows conductivity to be CONTROLLED by adding impurities (doping). Modern electronics rest entirely on band theory — research how this extends covalent bonding to industrial materials.

STRETCH

Research VSEPR theory (Valence Shell Electron Pair Repulsion) — Class 9 introduces simple Lewis structures, but VSEPR (Class 11+) predicts the 3D SHAPE of molecules. CH₄ is tetrahedral (109.5° angles); H₂O is bent (104.5° due to lone pair repulsion); NH₃ is pyramidal; CO₂ is linear. Molecular shape determines biological function — DNA, proteins, enzymes all work because of their 3D shapes. Research VSEPR theory and how it predicts molecular geometry.

Where else this chapter is tested

CBSE board isn't the only one — other exams test this chapter too.

AP Board SSC (Class 10) — ChemistryVery High — chemical bonding underlies Class 10 Periodic Table, Acids/Bases, Carbon Compounds, and Metallurgy

JEE Main and Advanced (Chemistry)Very High — chemical bonding (ionic, covalent, hybridisation, VSEPR) is among the most heavily tested JEE Chemistry topics

NEET (Chemistry)Very High — chemical bonding is core NEET Chemistry content

Chemistry Olympiad (INChO)High — bonding theory, electronegativity, molecular geometry, and intermolecular forces are tested with depth

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Questions students ask

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Electrical conduction requires CHARGED PARTICLES that can MOVE FREELY. In ionic compounds: ions ARE charged particles, but their mobility depends on state. SOLID STATE: ions are locked in a rigid crystal lattice — they vibrate in place but cannot flow. No movement = no current. MOLTEN STATE: heat breaks the lattice — ions can move freely throughout the liquid. They flow toward oppositely charged electrodes → current flows. DISSOLVED IN WATER (AQUEOUS): water dissociates the ionic compound — Na⁺ and Cl⁻ ions are surrounded by water molecules and can move freely → current flows. This is why your salt-water bath conducts (Na⁺ and Cl⁻ moving) but a salt crystal in your kitchen does not.

Calcium has valency 2 (it tends to lose 2 electrons to achieve argon configuration: Ca (2,8,8,2) − 2e⁻ → Ca²⁺ (2,8,8)). Chlorine has valency 1 (it tends to gain 1 electron: Cl (2,8,7) + 1e⁻ → Cl⁻ (2,8,8)). For the compound to be ELECTRICALLY NEUTRAL, total positive charges must equal total negative charges. One Ca²⁺ has +2 charge; to balance, we need 2 Cl⁻ ions (each −1, total −2). Result: CaCl₂ (one calcium, two chlorides). CHECK: +2 + (−1)×2 = 0 (correct). CaCl would have +2 + (−1) = +1 charge — not neutral, so cannot exist as a stable compound. The CROSS-MULTIPLY method automates this: Ca²⁺Cl¹⁻ → swap and write subscripts: Ca₁Cl₂ → CaCl₂.

The key is FORCES BETWEEN PARTICLES, not within them. IONIC COMPOUNDS: ions throughout the entire crystal are held by STRONG electrostatic attractions extending in 3D. To melt or boil, ALL these strong attractions must be overcome — requires very high temperatures (NaCl melts at 801°C). COVALENT COMPOUNDS: within each molecule (e.g., H₂O, CH₄), the covalent bonds are strong. BUT between molecules, only weak forces operate (van der Waals, hydrogen bonds). To melt or boil, only these WEAK intermolecular forces need to be overcome — requires low temperatures (CH₄ boils at −161°C; H₂O at 100°C). The covalent bonds within molecules are NOT broken on melting/boiling — only the weak forces between molecules. Result: covalent substances are usually gases or low-MP solids/liquids.

In a NORMAL COVALENT bond, each atom contributes ONE electron to the shared pair. For example, in H−H, each H contributes 1 electron, making the shared pair of 2 electrons. In a COORDINATE (DATIVE) bond, both electrons of the shared pair come from ONE atom (the donor). The other atom (acceptor) contributes 0 electrons. The donor must have a LONE PAIR (a pair of electrons not yet used in bonding). EXAMPLE: in ammonium ion NH₄⁺, the original NH₃ molecule has three N−H bonds (normal covalent) and one lone pair on N. When H⁺ (a proton with no electrons) approaches, the N donates BOTH electrons of its lone pair to form a coordinate bond with H⁺ → NH₄⁺. Once formed, the coordinate bond is identical in strength and properties to a normal covalent bond — only the way it formed is different.

Noble gases (He, Ne, Ar, Kr, Xe, Rn) already have STABLE outermost shell configurations: He has 2 (K shell full), Ne has 2,8 (L full), Ar has 2,8,8 (M with 8). They are at the THERMODYNAMIC MINIMUM of energy — they have no tendency to gain, lose, or share electrons because doing so would DESTABILISE them. Other elements try to ACHIEVE the noble gas configuration through bonding; noble gases ALREADY have it. This is why all elements that DO react try to look like the nearest noble gas: Na⁺ achieves Ne config, Mg²⁺ achieves Ne, Cl⁻ achieves Ar, etc. Note: under extreme conditions (high pressure, fluorine), heavier noble gases (Xe, Kr) can form unstable compounds (XeF₂, XeF₄, XeF₆) — but these were not discovered until 1962, after centuries of believing noble gases were completely inert.

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