Class 9 › Science › Chapter 9
Atomic Foundations of Matter
Laws of chemical combination, Dalton's atomic theory, covalent and ionic bonding, naming compounds, writing formulae, properties of ionic and covalent compounds, and molecular and formula unit mass.
1. Law of Conservation of Mass
Law of Conservation of Mass (Lavoisier, 1789)
Matter can neither be created nor destroyed in a chemical reaction: the total mass of reactants equals the total mass of products.
- Physical change (Activity 9.1): mass of salt solution = mass of water + salt; tearing paper does not change mass. Mass is the same in physical changes too.
- Antoine Lavoisier: Father of Modern Chemistry; 'in every operation an equal quantity of matter exists both before and after'.
Activity 9.2: vinegar + baking soda (sodium hydrogencarbonate) $\to$ carbon dioxide + other substances
Set-up 1 (balloon not attached; soda tipped in): readings differ because the CO2 gas escapes.
Set-up 2 (balloon fixed to flask mouth, closed system): the gas inflates the balloon and initial and final readings are the same.
A change in reading of ±1 in the last digit is experimental error. Keep both flask and balloon on the balance throughout.
Activity 9.3: sodium sulfate + barium chloride
Sodium sulfate + barium chloride $\to$ barium sulfate (white precipitate) + sodium chloride. Total mass is unchanged. (No gas forms, so an open system works.)
Solved examples
9.1: reactants $4.0 + 2.92 = 6.92$ g; products $1.76 + 0.72 + 4.44 = 6.92$ g. Equal, so the law is obeyed.
9.2: 12 g C + 32 g O $\to$ 44 g CO2. For 2.4 g C: $\dfrac{44}{12}\times 2.4 = 8.8$ g CO2.
- In gas-producing reactions use a closed system (or account for the gas); burning in an open beaker does not violate the law because the products escape as gases.
2. Law of Constant (Definite) Proportions
Proust's Law
In a compound, the elements are always combined in a fixed ratio by mass, irrespective of the source or method of preparation.
- Water from rivers, borewells or the ocean (purified) always has hydrogen : oxygen $= 1 : 8$ by mass; 9 g of water gives 1 g H and 8 g O.
- Cinnabar (hingula) always gives mercury 86.22% and sulfur 13.78%.
- Joseph Louis Proust: showed copper carbonate always has copper, carbon and oxygen in the same proportion by mass.
- Holds for compounds, not mixtures (a mixture's components can be in any ratio).
Example 9.3
NaCl: Na : Cl $= 23 : 35.5$. For 46 g Na: Cl needed $= \dfrac{35.5}{23}\times 46 = 71$ g.
Similar: CO has C : O $= 3:4$; 9 g C needs 12 g O. Ratios 4 : 1 and 8 : 2 are the same ratio.
3. Dalton's Atomic Theory
The two laws led to Dalton's atomic theory (1808), which explains why substances combine in fixed proportions and why mass is conserved. A postulate is a fundamental assumption accepted without formal proof.
Dalton's postulates
1. All matter is made of very tiny particles called atoms, which take part in chemical reactions.
2. Atoms are indivisible; they cannot be created or destroyed in a chemical reaction.
3. Atoms of a given element are identical in mass and chemical properties.
4. Atoms of different elements have different masses and chemical properties.
5. Atoms combine in the ratio of simple whole numbers to form compounds.
6. The relative number and kinds of atoms are constant in a given compound.
2. Atoms are indivisible; they cannot be created or destroyed in a chemical reaction.
3. Atoms of a given element are identical in mass and chemical properties.
4. Atoms of different elements have different masses and chemical properties.
5. Atoms combine in the ratio of simple whole numbers to form compounds.
6. The relative number and kinds of atoms are constant in a given compound.
- In a reaction atoms are rearranged, not destroyed: hydrogen + oxygen $\to$ water; magnesium burns to white magnesium oxide.
- Later corrections (Chapter 8): atoms are divisible (subatomic particles), and isotopes of an element differ in mass.
4. How Atoms Combine: Covalent Bonding
Molecule: an electrically neutral entity of more than one atom, capable of independent existence and showing all properties of that substance. Some elements (eg: helium) exist as single atoms because they are already stable.
- Atoms combine to achieve a stable octet (duplet for K shell) either by sharing electrons (covalent) or transferring them (ionic). The system's energy becomes lower, so it is more stable. The force holding atoms together is a chemical bond.
- Covalent bond: a shared pair of electrons between two atoms. One shared pair = single bond (—), two = double bond (=).
| Molecule | Valence electrons needed | Sharing | Bond |
|---|---|---|---|
| Hydrogen H2 | Each H needs 1 (K shell duplet) | Each shares 1 electron | H—H (single) |
| Chlorine Cl2 | Each Cl has 7, needs 1 | Each shares 1 | Cl—Cl (single) |
| Oxygen O2 | Each O has 6, needs 2 | Each shares 2 | O=O (double) |
| Hydrogen chloride HCl | H needs 1, Cl needs 1 | 1 pair shared | H—Cl (single) |
| Water H2O | O needs 2; each H needs 1 | O shares one electron with each of two H | H—O—H (two single) |
- Eg: nitrogen N2 (5 valence electrons, needs 3) shares 3 pairs: a triple bond; ammonia NH3, CO2 (C shares two pairs with each O), H2S follow the same method.
- Neon (2, 8) already has an octet: neither transfers nor shares electrons.
- Nuclear energy: splitting or combining nuclei releases enormous energy, used for electricity, medicine, research and space. Raja Ramanna is named in the chapter as contributing to India's nuclear energy programme.
Naming covalent compounds
- First element keeps its name; the second ends in -ide. Use prefixes: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6).
- Mono- is usually omitted for the first element but used for the second. If the prefix ends in 'o' or 'a' and the element starts with a vowel, drop the vowel (monoxide, pentoxide, tetroxide); prefixes ending in 'i' keep it (dioxide, trioxide).
- Hydrogen first: no prefix (H2S is hydrogen sulfide).
- Common names: H2O (water), NH3 (ammonia).
| Formula | Name | Formula | Name |
|---|---|---|---|
| CO | Carbon monoxide | PCl3 | Phosphorus trichloride |
| CO2 | Carbon dioxide | SF6 | Sulfur hexafluoride |
| CS2 | Carbon disulfide | N2O4 | Dinitrogen tetroxide |
| N2O5 | Dinitrogen pentoxide | H2S | Hydrogen sulfide |
5. Ionic Bonding
- Atoms with fewer than 4 valence electrons tend to donate them (become cations, positive); atoms with more than 4 tend to gain or share (gain makes anions, negative). Cations and anions together are ions.
- Ionic bond: the electrostatic force of attraction between oppositely charged ions, formed by transfer of electrons. Generally metals + non-metals.
Formation of sodium chloride by transfer of an electron.
- Crystal lattice: ionic compounds do not exist as single units; ions arrange in a repeating 3-D pattern. In NaCl each Na+ is surrounded by 6 Cl− and vice versa.
- Sulfur (6 valence electrons) gains 2 to form S2−; oxygen forms O2−; magnesium loses 2 to form Mg2+; MgCl2 has Mg2+ and two Cl−.
- Naming: cation first, then anion (simple anions end in -ide): sodium chloride, calcium oxide, magnesium sulfide. Polyatomic ions (groups of atoms) generally do not end in -ide.
| Cations | Anions |
|---|---|
| Na+, Li+, K+, Ag+, Cu+ (cuprous): 1+ Ca2+, Ba2+, Mg2+, Zn2+, Fe2+ (ferrous), Cu2+ (cupric): 2+ Al3+, Fe3+ (ferric): 3+ NH4+ (ammonium, polyatomic): 1+ | F−, Cl−, Br−, I−: 1− O2− (oxide), S2− (sulfide): 2− OH− (hydroxide), NO3− (nitrate), HCO3− (hydrogencarbonate): 1− CO32− (carbonate), SO42− (sulfate): 2− |
6. Writing Chemical Formulae
Criss-cross method
1. Write the symbols (cation first for ionic compounds).
2. Write the valencies / charges (numbers only) below.
3. Criss-cross the numbers as subscripts. A subscript of 1 is not written.
4. Divide by a common factor if there is one (simplest ratio). Put brackets around a polyatomic ion when two or more are needed.
| Compound | Ions / valencies | Criss-cross | Formula |
|---|---|---|---|
| Hydrogen chloride | H 1, Cl 1 | 1, 1 | HCl |
| Hydrogen sulfide | H 1, S 2 | 2, 1 | H2S |
| Carbon tetrachloride | C 4, Cl 1 | 1, 4 | CCl4 |
| Calcium chloride | Ca2+, Cl− | 1, 2 | CaCl2 |
| Aluminium oxide | Al3+, O2− | 2, 3 | Al2O3 |
| Magnesium oxide | Mg2+, O2− | 2, 2 → 1, 1 | MgO |
| Calcium carbonate | Ca2+, CO32− | 2, 2 → 1, 1 | CaCO3 |
| Magnesium hydroxide | Mg2+, OH− | 1, 2 | Mg(OH)2 |
| Aluminium hydroxide | Al3+, OH− | 1, 3 | Al(OH)3 |
| Aluminium sulfate | Al3+, SO42− | 2, 3 | Al2(SO4)3 |
- Charges are not written in the formula; the compound is neutral overall (total positive charge = total negative charge).
- Write Al(OH)3, not AlOH3. A formula of an ionic compound gives the simplest whole-number ratio of ions.
7. Properties of Ionic and Covalent Compounds
| Property | Ionic compounds | Covalent compounds |
|---|---|---|
| Examples | Sodium chloride, copper sulfate | Camphor, naphthalene, sugar, water |
| Formed by | Transfer of electrons (metal + non-metal) | Sharing of electrons (usually non-metals) |
| Solubility | Generally soluble in water; insoluble in kerosene and petrol | Most insoluble in water, soluble in kerosene and petrol (sugar is an exception: soluble in water) |
| Conductivity: solid | No (ions fixed in position) | No |
| Conductivity: in water | Yes (free-moving ions) | Generally no (sugar dissolves but gives no ions; camphor and naphthalene do not conduct) |
| Conductivity: molten | Yes, ions can move | No ions, so no |
| Melting and boiling points | High (strong inter-ionic attraction) | Usually low |
| Structure | 3-D crystal lattice; no molecules | Discrete molecules |
- Activity 9.4 safety: low-voltage battery for the conductivity test; petrol and kerosene are flammable.
- A solid that does not conduct but conducts when dissolved in water indicates an ionic compound.
8. Molecular Mass and Formula Unit Mass
Molecular mass
Covalent compounds
Sum of the atomic masses of all atoms in a molecule. Eg: H2O $= 2(1)+16 = 18$ u; CO2 $= 12 + 2(16) = 44$ u.
Formula unit mass
Ionic compounds
Ionic compounds do not form molecules; the formula unit is the simplest whole-number ratio of ions. Eg: Na2O $= 2(23)+16 = 62$ u.
Method with brackets: Ca(NO3)2 (Ca 40, N 14, O 16)
$40 + 2\times(14 + 3\times16) = 40 + 2\times62 = 164$ u. Multiply everything inside the bracket by the subscript.
Practice: HNO3 $= 1+14+48 = 63$ u; CH4 $= 12+4 = 16$ u; KCl $= 39+35.5 = 74.5$ u; Mg(OH)2 $= 24+2(16+1) = 58$ u.
- Unit: unified atomic mass unit (u).
9. Quick Revision
| Covalent bond | Ionic bond |
|---|---|
| Sharing of electron pair(s) | Transfer of electron(s): cation + anion |
| Forms molecules (H2, Cl2, O2, HCl, H2O) | Forms crystal lattice (NaCl, MgO) |
| Single, double, triple bonds | Electrostatic attraction between ions |
| Named with prefixes (di-, tri-…) | Cation name + anion name (-ide) |
| Molecular mass | Formula unit mass |
Laws
Two laws
Conservation of mass (Lavoisier 1789); constant proportions (Proust). Both support Dalton's theory (1808).
Tip
Formula check
Total + charge = total − charge. Use brackets for 2 or more polyatomic ions.
Tip
Valence electrons
Fewer than 4: lose (cation). 5 to 7: gain or share. 4: share. 8 (2 for He): stable.
Common mistakes:
• Saying mass is lost when a gas escapes (it escapes; in a closed system mass is the same).
• Applying the law of constant proportions to mixtures.
• Writing charges in formulae or forgetting brackets: write Al(OH)3, not AlOH3.
• Forgetting to simplify (Mg2O2 should be MgO).
• Calling ionic compounds 'molecules' (they have formula units).
• Saying solid ionic compounds conduct electricity (only molten or dissolved).
• Naming H2S 'dihydrogen sulfide' or CO 'carbon monooxide'.
• Giving oxygen two shared electrons in total instead of two each (O=O is a double bond).
• Saying mass is lost when a gas escapes (it escapes; in a closed system mass is the same).
• Applying the law of constant proportions to mixtures.
• Writing charges in formulae or forgetting brackets: write Al(OH)3, not AlOH3.
• Forgetting to simplify (Mg2O2 should be MgO).
• Calling ionic compounds 'molecules' (they have formula units).
• Saying solid ionic compounds conduct electricity (only molten or dissolved).
• Naming H2S 'dihydrogen sulfide' or CO 'carbon monooxide'.
• Giving oxygen two shared electrons in total instead of two each (O=O is a double bond).