Summary & Quick Quiz രാസബന്ധനം
The whole chapter on one page: the story of bonding in five pictures, the key boxes of the textbook, every rule you need – octet, ionic and covalent bonds, the 1.7 rule, valency and criss-cross formulae, the common ions and radicals – the mistakes to avoid, flashcards, a scored quiz and your study checklist.
1. The chapter in five pictures
2. The key boxes of the textbook
Having 8 electrons in the outermost shell, like the noble gases, is called the octet configuration. (Helium: 2 – the duplet.) – page 3
The force that holds the constituent particles together in a compound is the chemical bond. – page 3
The energy needed to remove the most loosely bound electron of the outermost shell of an isolated gaseous atom. – page 4
The energy released when an electron is added to a neutral gaseous atom to make it a negative ion. – page 4
In ionic compounds, the electrostatic force of attraction that holds the oppositely charged ions together is the ionic (electrovalent) bond. – page 4
Compounds formed by ionic bonds are ionic compounds or electrovalent compounds (NaCl, MgO). – page 5
The chemical bond formed by sharing electrons is the covalent bond. Sharing one pair gives a single bond: F–F, H–H. – page 6
Sharing two pairs → double bond (O=O); sharing three pairs → triple bond (N≡N). – page 6
Compounds formed by covalent bonds are covalent compounds. Non-metals combining usually give covalent compounds (HCl, H2O). – page 7
The relative ability of an atom in a covalent bond to attract the shared electron pair. Pauling scale 0–4; fluorine highest (3.98). – page 8
The number of electrons an atom gives, takes or shares when it takes part in a chemical reaction. – page 9
The total of the charges of the positive and negative ions in a salt is zero. – page 10
1875–1946, University of California: the electron pair, the covalent bond and electron dot structures; he also named the photon. – page 4
CO2, CCl4 and BeF2 are non-polar although their bonds have large electronegativity differences – because of their shape the pulls cancel. – page 8
H bonded to F, O or N gets δ+ and is attracted by an electronegative atom of the same or another molecule. It makes water a liquid and a universal solvent; ice is less dense than water. – page 8
Forces of attraction between molecules – hydrogen bonding is one of them. – page 8
Fe 2 (FeCl2, ferrous) and 3 (FeCl3, ferric); Cu 1 (Cu2O, cuprous) and 2 (CuO, cupric); P 3 (PCl3) and 5 (PCl5). – page 9
3. Facts & rules
📐 Facts & rules card
other atoms combine to reach an octet (hydrogen: a duplet, like helium)
metal + non-metal → transfer of electrons → ionic · non-metal + non-metal → sharing → covalent
Ionic or covalent? അയോണിക · സഹസംയോജക
| Ionic (electrovalent) compounds | Covalent compounds | |
|---|---|---|
| How formed | transfer of electrons: metal → cation (+), non-metal → anion (−); electrostatic attraction | sharing of electron pairs between atoms |
| Usually between | a metal and a non-metal (NaCl, MgO, CaCl2) | non-metals (H2, HCl, H2O, CH4, CO2) |
| Electronegativity difference | 1.7 or more | less than 1.7 |
| State | crystalline solids, hard, non-volatile | solids, liquids and gases |
| Melting / boiling point | very high | generally low |
| Solubility | soluble in polar solvents like water | generally insoluble in water; soluble in organic solvents (kerosene, carbon tetrachloride, benzene) |
| Electricity | does not conduct as a solid; conducts when molten or dissolved | generally does not conduct |
Single, double and triple bonds
| Shared pairs | Shared electrons | Bond | Examples |
|---|---|---|---|
| 1 | 2 | single bond ഏകബന്ധനം | H–H, F–F, Cl–Cl, H–Cl, H–F; H–O–H has two single bonds |
| 2 | 4 | double bond ദ്വിബന്ധനം | O=O; O=C=O |
| 3 | 6 | triple bond ത്രിബന്ധനം | N≡N |
The 1.7 rule ഇലക്ട്രോനെഗറ്റിവിറ്റി സ്കെയിൽ
Difference = larger value − smaller value (Fig 3.12). ≥ 1.7 → generally ionic; < 1.7 → covalent. In a covalent bond between different atoms the more electronegative atom gets δ− and the other δ+ (δ+H–Clδ−) – a polar bond. Equal atoms (H2, N2, O2) → non-polar.
The rule says “generally” – it is a guide. Formulae are shown for a metal with a non-metal.
Valency സംയോജകത
| Outermost electrons | The atom … | Valency | Examples |
|---|---|---|---|
| 1, 2, 3 | gives them away (metals) | 1, 2, 3 | Na 2, 8, 1 → 1 · Mg 2, 8, 2 → 2 · Al 2, 8, 3 → 3 |
| 4 | shares 4 | 4 | C 2, 4 → 4 (CCl4, CH4) |
| 5, 6, 7 | takes 3, 2, 1 (or shares them) | 8 − outermost | N 2, 5 → 3 · O 2, 6 → 2 · Cl 2, 8, 7 → 1 |
| 8 (He 2) | neither – stable | 0 | He, Ne, Ar |
Writing formulae – the criss-cross steps രാസസൂത്രം · പാദാങ്കം
- Write the symbols (or ions) – the metal / less electronegative one first: Al O.
- Write each valency (or charge number) above its symbol: Al 3, O 2.
- Swap them crosswise as subscripts: Al2O3.
- Divide by a common factor (Mg2O2 → MgO; C2O4 → CO2) and leave out 1.
- A radical taken more than once goes in brackets: Ca(OH)2, Mg3(PO4)2, (NH4)2SO4.
- Check: the total charge of a salt is zero.
Acids give H+ ions – their number is the basicity (HCl monobasic, H2SO4 dibasic, H3PO4 tribasic). Alkalis give OH− ions equal in number to the charge of the positive ion (NaOH, Ca(OH)2, Al(OH)3). Salts: acid + base → salt + water, e.g. NaOH + HCl → NaCl + H2O (നിർവീരീകരണ രാസപ്രവർത്തനം).
Common ions and radicals അയോൺ · റാഡിക്കൽ
| Positive ion | Name |
|---|---|
| Na+ | sodium |
| K+ | potassium |
| H+ | hydrogen |
| NH4+ | ammonium (radical) |
| Cu+ | cuprous – copper(I) |
| Mg2+ | magnesium |
| Ca2+ | calcium |
| Zn2+ | zinc |
| Fe2+ | ferrous – iron(II) |
| Cu2+ | cupric – copper(II) |
| Al3+ | aluminium |
| Fe3+ | ferric – iron(III) |
| Negative ion | Name |
|---|---|
| Cl− | chloride |
| F− | fluoride |
| OH− | hydroxide (radical) |
| NO3− | nitrate (radical) |
| O2− | oxide |
| S2− | sulphide |
| SO42− | sulphate (radical) |
| SO32− | sulphite (radical) |
| CO32− | carbonate (radical) |
| N3− | nitride |
| PO43− | phosphate (radical) |
- Changing the number of protons when an ion forms. Only electrons move: Na+ still has 11 protons, Cl− still 17.
- Writing Na− or Cl+. Metal atoms lose electrons → positive ions; non-metals gain → negative ions.
- Calling an ionic bond “sharing” or a covalent bond “transfer”. Ionic = transfer + attraction of ions; covalent = sharing of pairs.
- Counting shared electrons wrongly. O=O shares 2 pairs = 4 electrons; N≡N shares 3 pairs = 6 electrons.
- Drawing all the electrons in a dot diagram. Only the outermost electrons are drawn (Cl: 7, not 17).
- Taking the 1.7 rule the wrong way round, or subtracting the bigger value from the smaller. Always larger − smaller; 1.7 or more is ionic.
- Forgetting to simplify or to use brackets. Mg2O2 → MgO, C2O4 → CO2; write Mg3(PO4)2, not Mg3PO42.
- Writing the more electronegative element first. NaCl, not ClNa – the metal (less electronegative) comes first.
- Thinking every molecule with polar bonds is polar. CO2, CCl4, BeF2 are non-polar because of their shape.
- Saying ionic compounds conduct as solids. Only when molten or dissolved – then the ions can move.
4. Flashcards
Click a card to flip it.
5. Scored quiz
Answer all sixteen, then press Check my answers. Each wrong answer shows the page to revise. For number answers type only the number.
6. Your study checklist
Ticks come from “Mark this page as completed” at the bottom of each page.
Next: the model test paper ▶ – attempt it on paper with a timer first, then check with the solved paper.