Atomic Foundations of Matter explains what happens to atoms when substances react and combine. It starts with the Law of Conservation of Mass and the Law of Constant Proportions, shows how Dalton’s Atomic Theory explains both, and then covers covalent bonds (sharing electrons) and ionic bonds (transferring electrons). It ends with naming compounds, writing formulae by the criss-cross method, comparing ionic and covalent compounds, and calculating molecular mass and formula unit mass. These skills are used in every later chemistry chapter.
Key Concepts
1. The Law of Conservation of Mass
Hydrogen and oxygen are gases; hydrogen burns and oxygen supports burning. The water they form is a liquid that puts out fire. The properties change completely, yet the mass of water formed equals the sum of the masses of hydrogen and oxygen that combined. The chapter tests this for a physical change and then for chemical changes.
Physical change (Activity 9.1): a beaker with about 50 mL of water and a spatula of common salt is weighed on a digital balance, then swirled until the salt dissolves and weighed again. The readings match: the mass of the solution equals the mass of water plus salt. There is practically no change in mass in a physical change. Tearing paper into pieces gives the same result.
Chemical change (Activity 9.2)
Vinegar + Baking soda (sodium hydrogencarbonate) → Carbon dioxide + Other substances
| Feature | Set-up 1 (open) | Set-up 2 (closed) |
|---|---|---|
| Arrangement | Flask with about 20 mL vinegar; balloon with about 2 g baking soda kept beside it on the balance | Balloon with the baking soda tied over the mouth of the flask with a thread |
| Observation | Brisk effervescence | Brisk effervescence; the balloon inflates |
| Final reading | Lower than the initial reading | Same as the initial reading |
| Reason | Carbon dioxide escapes and is not weighed | Carbon dioxide is trapped in the balloon and weighed |
Set-up 2 shows that the total mass of vinegar and baking soda equals the total mass of carbon dioxide and the other substances formed. Any measurement has an uncertainty of ±1 in the last digit, so a change only in the last digit of the balance reading is within experimental error.
Statement of the law
Matter can neither be created nor destroyed in a chemical reaction. This is the Law of Conservation of Mass, proposed by Antoine Lavoisier in 1789, known as the Father of Modern Chemistry. In practice: total mass of reactants = total mass of products.
Verification with a precipitate (Activity 9.3)
About 10 mL of 1% m/v sodium sulfate solution (flask A) and 10 mL of 1% m/v barium chloride solution (flask B) are weighed together. B is poured into A. A white precipitate forms:
Sodium sulfate + Barium chloride → Barium sulfate + Sodium chloride
Both flasks are weighed again and the reading is unchanged. An open system works here because no gas is formed.
Open and closed systems: if a gaseous product escapes, the mass left behind falls. In a closed container the reading stays the same.
2. The Law of Constant Proportions
Joseph Proust found that in any compound formed by two or more elements, the elements combine in a fixed ratio by mass, whatever the source. He showed that copper carbonate always contains copper, carbon and oxygen in the same proportion by mass, however it was prepared. This is the Law of Constant Proportions, also called the Law of Definite Proportions or Proust’s Law.
Water: purified water from a river, a borewell or the ocean always contains hydrogen and oxygen in a mass ratio of 1 : 8. Decomposing 9 g of it always gives 1 g of hydrogen and 8 g of oxygen. So 18 g of water contains 2 g of hydrogen and 16 g of oxygen.
Cinnabar: the red pigment called hingula in India and cinnabar in Latin and English gives mercury and sulfur on heating, in mass percentages of around 86.22% and 13.78%.
Calculations: sodium chloride has sodium and chlorine in the mass ratio 23 : 35.5. For 46 g of sodium, chlorine = (35.5 ÷ 23) × 46 = 71 g.
Compounds and mixtures: the law holds for compounds only. A mixture such as salt in water can be made in any proportion, so it has no fixed composition. Ratios such as 4 : 1 and 8 : 2 are the same ratio and both obey the law.
3. Dalton’s Atomic Theory
John Dalton, who taught in Manchester from 1793, presented his atomic theory in 1808 as a set of postulates. A postulate is a fundamental assumption accepted as truth without formal proof, from which further ideas are formed or developed.
- All matter is made up of very tiny particles called atoms, which participate in chemical reactions.
- Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction.
- Atoms of a given element are identical in mass and chemical properties.
- Atoms of different elements have different masses and chemical properties.
- Atoms combine in the ratio of simple whole numbers to form compounds.
- The relative number and kinds of atoms are constant in a given compound.
| Law | Postulates | Reasoning |
|---|---|---|
| Conservation of Mass | 2 and 3 | In a reaction the same atoms are only rearranged, each with its fixed mass, so the total mass cannot change. |
| Constant Proportions | 3, 5 and 6 | A compound always has the same kinds of atoms in the same relative numbers, each with a fixed mass, so the mass ratio is fixed. |
When magnesium burns in air, a white powder of magnesium oxide forms: magnesium atoms have combined with oxygen atoms, and no atom is destroyed. You learnt in Chapter 8 that atoms contain electrons, protons and neutrons. For this chapter the key point still holds: atoms are neither created nor destroyed in a chemical reaction.
4. Molecules and Covalent Bonds (Sharing of Electrons)
A molecule is an electrically neutral entity consisting of more than one atom that is capable of independent existence and shows all the properties of that substance. A hydrogen molecule has two hydrogen atoms; a hydrogen chloride molecule has one hydrogen and one chlorine atom. Some elements, such as helium, exist only as atoms because their atoms are already stable.
Why atoms combine: an atom with 8 electrons in its valence shell is stable, and so is one whose only shell (K) holds 2 electrons. Atoms with incomplete valence shells share, gain or lose electrons. When atoms combine, the total energy becomes lower than that of the separate atoms, so the arrangement is more stable. The force that holds atoms together is a chemical bond. Sharing electrons gives a covalent bond; transferring electrons gives an ionic bond.
Molecules of elements
H2: hydrogen has one electron in its K-shell and needs one more. Two hydrogen atoms share one electron each. The shared pair attracts both nuclei and holds them together. Interaction through a shared pair of electrons is a covalent bond; one shared pair is a single bond, H-H.
Cl2: chlorine (2, 8, 7) needs one electron. Two chlorine atoms share one electron each: Cl-Cl, a single bond. Fluorine (2, 7) forms F2 the same way.
O2: oxygen (2, 6) needs two electrons. Two oxygen atoms share two electrons each, so two pairs are shared: a double bond, O=O.
N2: nitrogen (2, 5) needs three electrons. The two atoms share three electrons each, so three pairs are shared: a triple bond, N≡N.
Molecules of compounds
Write both configurations, then count what each atom needs. In HCl, hydrogen needs one electron for a duplet and chlorine one for an octet, so they share one electron each: H-Cl. Bonded by sharing, HCl is a covalent compound. In H2O, oxygen needs two electrons and each hydrogen one, so two hydrogen atoms each share an electron with one oxygen atom: H-O-H.
| Molecule | Electrons shared | Bonds |
|---|---|---|
| CO2 | C shares 4, each O shares 2 | Two double bonds, O=C=O |
| NH3 | N shares 3, each H shares 1 | Three single N-H bonds |
| H2S | S shares 2, each H shares 1 | Two single bonds, H-S-H |
Neon (2, 8) already has a complete octet, so it neither transfers nor shares electrons.
Naming covalent compounds
Prefixes give the number of atoms: mono (1), di (2), tri (3), tetra (4), penta (5), hexa (6). The first element keeps its name and the second ends in -ide. Mono- is usually left out for the first element but used for the second. If a prefix ends in ‘o’ or ‘a’ and the element starts with a vowel, drop that vowel (monoxide, pentoxide); a prefix ending in ‘i’ is kept (dioxide, trioxide). Hydrogen as the first element takes no prefix. H2O and NH3 keep their common names, water and ammonia.
| Formula | Name | Formula | Name |
|---|---|---|---|
| CO | carbon monoxide | SF6 | sulfur hexafluoride |
| CO2 | carbon dioxide | N2O4 | dinitrogen tetroxide |
| CS2 | carbon disulfide | N2O5 | dinitrogen pentoxide |
| PCl3 | phosphorus trichloride | H2S | hydrogen sulfide |
5. Ions and Ionic Bonds (Transfer of Electrons)
An atom with fewer than four valence electrons generally gives them away to become stable. Atoms with more than four usually gain or share electrons.
Cation: sodium (11; 2, 8, 1) loses its one valence electron. With 11 protons and 10 electrons it carries one positive charge: the sodium cation. Na → Na+ + e−
Anion: chlorine (2, 8, 7) gains one electron. With 17 protons and 18 electrons it carries one negative charge: the chloride anion. Cl + e− → Cl−
Bond: cations and anions are together called ions. The electrostatic force of attraction between oppositely charged ions that holds them together is the ionic bond. Na + Cl → Na+ + Cl− → NaCl
Crystals: ionic compounds form 3-D crystals. In sodium chloride each Na+ is surrounded by six Cl− ions and each Cl− by six Na+ ions. This repeating arrangement is the crystal structure; drawn with ions as points it is a crystal lattice. Ionic compounds therefore do not form molecules.
Larger charges: sulfur (six outer electrons) gains two to become S2−; oxygen forms O2− the same way. Magnesium gives two electrons to become Mg2+; each chlorine takes one, so one Mg2+ combines with two Cl− to give MgCl2. In sodium sulfide, two sodium atoms each give one electron to one sulfur atom: Na2S.
Naming ionic compounds
Name the cation first, then the anion. Simple anions end in -ide. Metals generally form cations and non-metals anions, so ionic compounds typically form from a metal and a non-metal: sodium chloride, calcium oxide, magnesium sulfide. Ions made of atoms of two or more elements are polyatomic ions; their names generally do not end in -ide.
| Valency | Cations | Anions |
|---|---|---|
| 1 | Na+, Li+, K+, Ag+, cuprous Cu+, ammonium NH4+ | F−, Cl−, Br−, I−, hydroxide OH−, nitrate NO3−, hydrogencarbonate HCO3− |
| 2 | Ca2+, Ba2+, ferrous Fe2+, cupric Cu2+, Mg2+, Zn2+ | oxide O2−, sulfide S2−, carbonate CO32−, sulfate SO42− |
| 3 | ferric Fe3+, Al3+ | - |
“-ous” goes with the lower charge and “-ic” with the higher: ferrous Fe2+, ferric Fe3+, cuprous Cu+, cupric Cu2+.
6. Properties of Ionic and Covalent Compounds
Activity 9.4 tests camphor, sodium chloride, copper sulfate, sugar and naphthalene for solubility in water, kerosene and petrol, and for electrical conductivity. Two carbon or metal electrodes pushed through a piece of cardboard are connected to a 9 V battery and a bulb; a glowing bulb means the sample conducts. Do not touch the electrodes, and keep kerosene and petrol away from flames.
- Ionic compounds do not conduct as solids because their ions are held in fixed positions by strong forces. Dissolved in water, the ions are free to move and the compound conducts. Melting also frees the ions, so a molten ionic compound conducts.
- Sugar dissolves in water but gives no ions, so its solution does not conduct. Camphor and naphthalene also do not conduct: covalent compounds are made of neutral molecules.
- Strong attraction between ions gives ionic compounds generally high melting and boiling points; covalent compounds usually have low ones.
| Property | Ionic compounds | Covalent compounds |
|---|---|---|
| How formed | Transfer of electrons | Sharing of electrons |
| Particles | Cations and anions in a 3-D crystal | Separate neutral molecules |
| Solubility in water | Generally soluble | Mostly insoluble (sugar is an exception) |
| Solubility in kerosene, petrol | Insoluble | Mostly soluble |
| Conductivity as a solid | Do not conduct | Do not conduct |
| Conductivity in water solution | Conduct | Do not conduct |
| Melting and boiling points | Generally high | Usually low |
| Examples | Sodium chloride, copper sulfate | Camphor, naphthalene, sugar |
7. Writing Chemical Formulae
The quick route to a formula is the criss-cross method using valencies.
Covalent compounds
- Write the symbols of the elements.
- Write the valency of each element below its symbol.
- Cross over the valencies and write them as subscripts after the symbols. A subscript of 1 is not written.
| Compound | Symbols | Valencies | Formula |
|---|---|---|---|
| Hydrogen chloride | H, Cl | 1, 1 | HCl |
| Hydrogen sulfide | H, S | 1, 2 | H2S |
| Carbon tetrachloride | C, Cl | 4, 1 | CCl4 |
Ionic compounds
- Write the symbol of the cation first and the anion second.
- Write the charges under the symbols.
- Cross over only the numbers of the charges, not the signs.
- The formula shows the simplest ratio, so divide the subscripts by any common factor. Subscripts 2 and 4 become 1 and 2.
The charges on the ions are not shown in the final formula. The total positive charge must equal the total negative charge, so the compound is neutral.
| Compound | Ions and charges | After criss-cross | Final formula |
|---|---|---|---|
| Calcium chloride | Ca 2+, Cl 1− | Ca1Cl2 | CaCl2 |
| Aluminium oxide | Al 3+, O 2− | Al2O3 | Al2O3 |
| Magnesium oxide | Mg 2+, O 2− | Mg2O2 | MgO (divide by 2) |
| Calcium carbonate | Ca 2+, CO3 2− | Ca2(CO3)2 | CaCO3 (divide by 2) |
| Magnesium hydroxide | Mg 2+, OH 1− | Mg1(OH)2 | Mg(OH)2 |
| Aluminium hydroxide | Al 3+, OH 1− | Al1(OH)3 | Al(OH)3 |
| Aluminium sulfate | Al 3+, SO4 2− | Al2(SO4)3 | Al2(SO4)3 |
Brackets for polyatomic ions
Use brackets when two or more polyatomic ions of the same type appear in a formula. In Al(OH)3 the bracket with subscript 3 means three hydroxide ions are bound to one aluminium ion. Writing AlOH3 would wrongly mean one oxygen and three hydrogen atoms. Brackets are not needed when only one polyatomic ion is present: NaOH, CaCO3, NH4Cl.
What a formula tells you
- The formula of a covalent compound shows the elements and the number of atoms of each element present in one molecule.
- The formula of an ionic compound shows the simplest whole number ratio of the atoms of the different elements present.
8. Molecular Mass and Formula Unit Mass
Once the formula is known, the mass of one unit of the substance is found by adding up the masses of the atoms in it. Atomic masses are in atomic mass units, u.
Molecular mass (covalent compounds)
Molecular mass is the total mass of a molecule, calculated by adding the atomic masses of all the atoms in it.
- H2O: (1 u × 2) + (16 u × 1) = 18 u
- CO2: (12 u × 1) + (16 u × 2) = 44 u
- HNO3: (1 u × 1) + (14 u × 1) + (16 u × 3) = 1 + 14 + 48 = 63 u
- CH4: (12 u × 1) + (1 u × 4) = 16 u
Formula unit mass (ionic compounds)
Ionic compounds form 3-D crystals and have no separate molecules, so the term molecular mass does not suit them. In an ionic compound, the collection of the simplest whole number ratio of ions is called a formula unit. Its mass is the formula unit mass, the sum of the atomic masses of all the atoms in one formula unit.
- Na2O: (23 u × 2) + (16 u × 1) = 62 u
- Ca(NO3)2: (40 u × 1) + {(14 u × 1) + (16 u × 3)} × 2 = 40 + 62 × 2 = 164 u
- KCl: 39 u + 35.5 u = 74.5 u
- Mg(OH)2: 24 u + (16 u + 1 u) × 2 = 24 + 34 = 58 u
Method for brackets: find the mass of the bracketed group, then multiply by the subscript outside.
Formula and Reaction Sheet
| Item | Formula, rule or reaction | When to use it |
|---|---|---|
| Law of Conservation of Mass | Total mass of reactants = total mass of products | Verification sums; finding a missing product mass |
| Law of Constant Proportions | Mass of B = (ratio part of B ÷ ratio part of A) × mass of A | Fixed-ratio sums on a compound |
| Activity 9.2 reaction | Vinegar + Baking soda (sodium hydrogencarbonate) → Carbon dioxide + Other substances | Open versus closed system |
| Activity 9.3 reaction | Sodium sulfate + Barium chloride → Barium sulfate (white precipitate) + Sodium chloride | Verifying the law with no gas |
| Cation formation | Na → Na+ + e−; Mg → Mg2+ + 2e− | Metals with 1 to 3 valence electrons |
| Anion formation | Cl + e− → Cl−; S + 2e− → S2− | Non-metals with 5 to 7 valence electrons |
| Criss-cross (ionic) | Cation with charge a+, anion with charge b− gives CationbAniona, then simplify | Any metal with any anion |
| Molecular mass | Sum of (atomic mass × number of atoms) over all elements | Covalent molecules: H2O = 18 u, CO2 = 44 u |
| Formula unit mass | Sum of atomic masses in one formula unit | Ionic compounds: Na2O = 62 u, Ca(NO3)2 = 164 u |
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Important Definitions
| Term | Definition |
|---|---|
| Law of Conservation of Mass | Matter (mass) can neither be created nor destroyed in a chemical reaction. Proposed by Antoine Lavoisier in 1789. |
| Law of Constant Proportions | In any compound formed by two or more elements, the elements combine in a fixed ratio by mass, irrespective of the source. Also called the Law of Definite Proportions or Proust’s Law. |
| Postulate | A fundamental assumption accepted as truth without formal proof, from which further ideas are formed or developed. |
| Molecule | An electrically neutral entity consisting of more than one atom that is capable of independent existence and shows all the properties of that substance. |
| Chemical bond | The force that holds atoms together. |
| Covalent bond | The interaction between atoms through a shared pair of electrons. |
| Cation | A positively charged ion formed when an atom loses one or more electrons, such as Na+. |
| Anion | A negatively charged ion formed when an atom gains one or more electrons, such as Cl−. |
| Ions | Cations and anions collectively. |
| Ionic bond | The electrostatic force of attraction between oppositely charged ions that holds them together. |
| Polyatomic ion | An ion formed by the combination of atoms of two or more elements, such as SO42− and NH4+. |
| Crystal structure | The regular, repeating 3-D arrangement of oppositely charged ions in an ionic solid. Drawn with ions as points, it is a crystal lattice. |
| Molecular mass | The total mass of a molecule, calculated by adding the atomic masses of all the atoms in it. |
| Formula unit | In an ionic compound, the collection of the simplest whole number ratio of ions. |
| Formula unit mass | The mass of a formula unit: the sum of the atomic masses of all the atoms in it. |
Solved Examples (NCERT-Based)
Example 1: Verifying conservation of mass (NCERT Example 9.1)
In a group activity, students place 4.0 g of calcium carbonate with 2.92 g of hydrochloric acid in a closed container. After the reaction is over, they measured 1.76 g of carbon dioxide, 0.72 g of water, and 4.44 g of calcium chloride. Verify whether the Law of Conservation of Mass is obeyed or not.
Solution: Total mass of reactants = 4.0 g + 2.92 g = 6.92 g.
Total mass of products = 1.76 g + 0.72 g + 4.44 g = 6.92 g.
Mass of reactants = mass of products, so the Law of Conservation of Mass is obeyed.
Example 2: A second verification (NCERT Exercise 11)
5.3 g of sodium carbonate and 6.0 g of acetic acid react to produce 2.2 g of carbon dioxide, 0.9 g of water, and 8.2 g of sodium acetate. Verify whether the law of conservation of mass is valid.
Solution: Total mass of reactants = 5.3 g + 6.0 g = 11.3 g.
Total mass of products = 2.2 g + 0.9 g + 8.2 g = 11.3 g.
The two totals are equal, so the law is valid for this reaction.
Example 3: Scaling a product mass (NCERT Example 9.2)
12 g of carbon combines with 32 g of oxygen to form 44 g of carbon dioxide as per the given equation: Carbon + Oxygen → Carbon dioxide. If 2.4 g of carbon reacts completely with oxygen, how much carbon dioxide will be produced?
Solution: 12 g of carbon gives 44 g of carbon dioxide.
So 1 g of carbon gives 44/12 g of carbon dioxide.
2.4 g of carbon gives (44/12) × 2.4 g = 44 × 0.2 g = 8.8 g of carbon dioxide.
Example 4: Using a fixed mass ratio (NCERT Example 9.3)
Sodium chloride (NaCl) contains sodium and chlorine in the mass ratio of 23:35.5. If 46 g of sodium reacts completely, how much chlorine is needed to form NaCl?
Solution: Mass of chlorine required = (35.5 ÷ 23) × 46 = 35.5 × 2 = 71 g.
The mass of NaCl formed is 46 + 71 = 117 g, by the Law of Conservation of Mass.
Example 5: Percentage composition and hydrogen with oxygen (Pause and Ponder 2, 3 and 4)
(a) When 20 g of hydrogen reacts completely with 160 g of oxygen, how much water is formed according to the Law of Conservation of Mass? (b) A compound consists of 40% sulfur and 60% oxygen by mass. In a sample of the same compound containing 20 g of sulfur, what mass of oxygen must be present to satisfy the Law of Constant Proportions? (c) Carbon monoxide (CO) contains carbon and oxygen in the mass ratio of 3:4. How much oxygen will combine with 9 g of carbon to form carbon monoxide?
Solution (a): Mass of water = 20 g + 160 g = 180 g. The ratio 20 : 160 = 1 : 8 also matches the fixed composition of water.
Solution (b): Sulfur : oxygen = 40 : 60 = 2 : 3. Oxygen = (60 ÷ 40) × 20 g = 30 g.
Solution (c): Oxygen = (4 ÷ 3) × 9 g = 12 g.
Example 6: Predicting ions, bond and formula (NCERT Exercise 1)
A particular element (A) has one electron in its third shell. There is another element (B) with six electrons in its second shell. (i) How many electrons does A tend to give or take to become stable? (ii) What kind of ion would it form? (iii) How many electrons does B tend to give or take to become stable? (iv) What kind of ion would it form? (v) If A and B were to combine, what kind of bond would be formed? (vi) What would be the formula for the compound thus formed?
Solution: A has one electron in its third shell, so its configuration is 2, 8, 1 (sodium, atomic number 11). B has six electrons in its second shell, so its configuration is 2, 6 (oxygen, atomic number 8).
(i) A tends to give 1 electron. (ii) It forms a cation with charge 1+, A+.
(iii) B tends to take 2 electrons. (iv) It forms an anion with charge 2−, B2−.
(v) Electrons are transferred from A to B, so an ionic bond forms.
(vi) Criss-cross the charges 1 and 2: the formula is A2B. Two A atoms each give one electron to one B atom. With the real elements this is Na2O, sodium oxide.
Example 7: Formulae from pairs of ions (NCERT Exercise 6 and Pause and Ponder 18)
Write the formulae of the compounds formed from the following pairs of ions: (i) Ca2+ and Br− (ii) Al3+ and CO32− (iii) K+ and SO42− (iv) NH4+ and Cl− (v) Fe3+ and OH− (vi) K+ and CO32−
Solution: Criss-cross the charge numbers and simplify.
| Pair | Charges | Formula |
|---|---|---|
| Ca2+, Br− | 2, 1 | CaBr2 |
| Al3+, CO32− | 3, 2 | Al2(CO3)3 |
| K+, SO42− | 1, 2 | K2SO4 |
| NH4+, Cl− | 1, 1 | NH4Cl |
| Fe3+, OH− | 3, 1 | Fe(OH)3 |
| K+, CO32− | 1, 2 | K2CO3 |
Example 8: Formula unit masses (NCERT Exercise 8)
Determine the formula unit mass of the following substances. (i) Ammonium nitrate (NH4NO3), used as a nitrogen fertiliser, which is essential for plant growth. (ii) Phosphoric acid (H3PO4), used to make phosphate fertiliser and detergents. (iii) Sodium hydrogencarbonate (NaHCO3), used to relieve acidity and helps in digestion.
Solution: Atomic masses: H = 1 u, C = 12 u, N = 14 u, O = 16 u, Na = 23 u, and P = 31 u (the standard value; the chapter does not list phosphorus).
(i) NH4NO3: N appears twice, H four times, O three times. (14 × 2) + (1 × 4) + (16 × 3) = 28 + 4 + 48 = 80 u.
(ii) H3PO4: (1 × 3) + (31 × 1) + (16 × 4) = 3 + 31 + 64 = 98 u.
(iii) NaHCO3: 23 + 1 + 12 + (16 × 3) = 23 + 1 + 12 + 48 = 84 u.
Example 9: Identifying a species from its particles (NCERT Exercise 12)
If a species has 11 protons, 12 neutrons and 10 electrons then (i) what is its atomic number and mass number? (ii) is it neutral, a cation or an anion? Explain. (iii) write its electronic configuration. (iv) name the species.
Solution: (i) Atomic number = number of protons = 11. Mass number = protons + neutrons = 11 + 12 = 23.
(ii) It has 11 positive protons and only 10 negative electrons, one electron short. It carries a charge of +1, so it is a cation.
(iii) Electronic configuration of the 10 electrons: 2, 8.
(iv) Atomic number 11 is sodium, so the species is the sodium ion, Na+.
Competency-Based Questions (with answers)
1. Case-based: Burning in an open beaker
A student burns 10 g of ethanol in an open beaker. After the reaction, no residue is left in the beaker. A friend says the Law of Conservation of Mass has failed. A second student burns a strip of magnesium ribbon in air and finds that the white powder collected weighs more than the ribbon did.
(a) Is the law violated in the ethanol case? No. Burning ethanol produces gases (carbon dioxide and water vapour) that escape from the open beaker. The mass of ethanol plus the oxygen used equals the mass of these gases.
(b) Why does the magnesium product weigh more? The white powder is magnesium oxide. Oxygen from the air combined with the magnesium, and its mass is now part of the product.
(c) How would you redesign either experiment to verify the law? Carry out the reaction in a closed system on the balance, as in set-up 2 of Activity 9.2, so that no gas enters or leaves and the balance reading can be compared before and after.
2. Case-based: Water from many sources
Samples of water are collected from a river, a borewell and the ocean. Each is purified and decomposed completely.
(a) Which law predicts that all three samples give the same result? The Law of Constant Proportions (Law of Definite Proportions, Proust’s Law).
(b) A purified sample of 36 g is decomposed. What masses of hydrogen and oxygen are obtained? Hydrogen : oxygen = 1 : 8, so 1 part in 9 is hydrogen. Hydrogen = 36 × 1/9 = 4 g. Oxygen = 36 × 8/9 = 32 g. Check: 4 + 32 = 36 g.
(c) Why must the water be purified first? Unpurified water is a mixture containing dissolved salts and other substances in varying amounts. Only the pure compound has a fixed composition.
3. Assertion-Reason (NCERT Exercise 14)
Assertion (A): Copper sulfate conducts electricity in the molten state but not in the solid state. Reason (R): Copper and sulfate ions are fixed in the lattice in molten state, while in solid state they can move freely.
(i) Both A and R are true, and R is the correct explanation of A. (ii) Both A and R are true, but R is not the correct explanation of A. (iii) A is true, but R is false. (iv) A is false, but R is true.
Answer: (iii) A is true, but R is false. The reason has the two states swapped. In the solid the ions are fixed in the lattice; in the molten state they are free to move, and moving ions carry the current.
4. Assertion-Reason
Assertion (A): A solution of sugar in water does not conduct electricity. Reason (R): Sugar is a covalent compound that does not provide ions in solution.
(i) Both A and R are true, and R is the correct explanation of A. (ii) Both A and R are true, but R is not the correct explanation of A. (iii) A is true, but R is false. (iv) A is false, but R is true.
Answer: (i) Both A and R are true, and R is the correct explanation of A. Conduction in a solution needs free ions. Sugar dissolves as neutral molecules, so there is nothing to carry charge.
5. Assertion-Reason
Assertion (A): Helium exists only as single atoms and does not form He2 molecules. Reason (R): A helium atom has two electrons in its K-shell, which is its only shell, so its valence shell is complete and stable.
(i) Both A and R are true, and R is the correct explanation of A. (ii) Both A and R are true, but R is not the correct explanation of A. (iii) A is true, but R is false. (iv) A is false, but R is true.
Answer: (i) Both A and R are true, and R is the correct explanation of A. A K-shell holding two electrons is complete and stable, so helium has no need to share, gain or lose electrons.
6. Source-based: Reading Activity 9.4 results
A student tests three unlabelled white solids, P, Q and R, and records: P dissolves in water, not in kerosene, and its water solution lights the bulb. Q dissolves in kerosene, not in water, and does not light the bulb. R dissolves in water, not in kerosene, and its solution does not light the bulb. None of the solids lights the bulb.
(a) Which solid is ionic? P. It dissolves in water, is insoluble in kerosene, and conducts only when dissolved, because its ions become free to move.
(b) Which of the chapter’s compounds could Q and R be? Q behaves like camphor or naphthalene (a covalent compound soluble in kerosene). R behaves like sugar (covalent, soluble in water, no ions in solution).
(c) Why does none of the solids conduct? In P the ions are held in fixed positions in the crystal. Q and R contain no ions at all.
7. Case-based: Designing a neutral ionic compound (NCERT Exercise 3)
You want to design a new ionic compound, where the total positive charge is 6+ and the total negative charge is 6−. Which of the following combinations gives the correct number of ions? (i) 2 Al3+ and 3 Cl− (ii) 3 Mg2+ and 1 PO43− (iii) 2 Fe3+ and 3 O2− (iv) 3 Ca2+ and 2 SO42−
Answer: (iii). Work out both totals for each option.
| Option | Total positive | Total negative | Balanced for 6+ and 6− |
|---|---|---|---|
| (i) | 2 × 3 = 6+ | 3 × 1 = 3− | No |
| (ii) | 3 × 2 = 6+ | 1 × 3 = 3− | No |
| (iii) | 2 × 3 = 6+ | 3 × 2 = 6− | Yes: Fe2O3 |
| (iv) | 3 × 2 = 6+ | 2 × 2 = 4− | No |
8. Case-based: The unknown metal M (Pause and Ponder 20)
Metal M, with two electrons in its valence shell (M shell), reacts with oxygen to form a compound that is slightly soluble in water. Predict its: (i) formula (ii) type of bond (iii) electrical conductivity of its aqueous solution.
Answer: The M shell is the third shell, so M has configuration 2, 8, 2 (magnesium). It loses two electrons to form M2+, and oxygen gains two to form O2−.
(i) Criss-cross gives M2O2, which simplifies to MO. (ii) Electron transfer gives an ionic bond. (iii) The part that dissolves releases ions, so its aqueous solution conducts electricity.
Important Questions for Exams
1-Mark Questions
- State the Law of Conservation of Mass. Mass can neither be created nor destroyed in a chemical reaction; the total mass of reactants equals the total mass of products.
- Who proposed the Law of Constant Proportions? Joseph Proust.
- What is a formula unit? In an ionic compound, the collection of the simplest whole number ratio of ions.
- Name N2O5 and CS2. Dinitrogen pentoxide and carbon disulfide.
- What kind of ion will oxygen (O) form? An anion, the oxide ion O2−, by gaining two electrons.
2-Mark Questions
- The Law of Definite Proportions holds true for compounds but not for mixtures. Give reason. In a compound the elements combine in fixed numbers of atoms, so the mass ratio is fixed whatever the source. A mixture can be made with its components in any proportion, so it has no fixed composition.
- Neon (atomic number 10) neither transfers nor shares its valence electrons. Explain. Neon’s configuration is 2, 8. Its valence shell already has a complete octet, so it is stable and has no tendency to lose, gain or share electrons.
- Why is the formula of aluminium hydroxide written as Al(OH)3 and not AlOH3? The formula has three hydroxide ions for each aluminium ion. The bracket shows that the subscript 3 applies to the whole OH group. AlOH3 would mean one oxygen and three hydrogen atoms.
- Find the formula unit mass of KCl and Mg(OH)2. (K = 39 u, Cl = 35.5 u, Mg = 24 u, O = 16 u, H = 1 u) KCl = 39 + 35.5 = 74.5 u. Mg(OH)2 = 24 + (16 + 1) × 2 = 58 u.
- Write the formulae: (i) Sodium hydrogencarbonate (ii) Sulfur dioxide (iii) Ferric chloride (iv) Cuprous oxide. (i) NaHCO3 (ii) SO2 (iii) FeCl3 (iv) Cu2O.
3-Mark Questions
- Explain the formation of sodium chloride from sodium and chlorine atoms. Sodium (2, 8, 1) loses its one valence electron to become Na+ (11 protons, 10 electrons): Na → Na+ + e−. Chlorine (2, 8, 7) gains that electron to become Cl− (17 protons, 18 electrons): Cl + e− → Cl−. The oppositely charged ions are held by electrostatic attraction, the ionic bond, forming NaCl.
- Choose the correct statement(s) and correct the false statement(s). (i) Elements are made up of molecules and compounds are made up of atoms. (ii) The molecule of a compound is always made up of two or more atoms of the same kind. (iii) One molecule of nitrogen gas contains three nitrogen atoms. (iv) Water is made of two hydrogen atoms, covalently bonded with one oxygen atom. (i) False. Elements are made of atoms of one kind (some as single atoms, like helium, some as molecules, like H2); compounds are made of atoms of two or more different elements combined. (ii) False. A molecule of a compound is made of atoms of two or more different kinds. (iii) False. A nitrogen molecule, N2, contains two nitrogen atoms. (iv) True.
- Two elements, A and B, have the following configurations: A: 2, 8, 5 and B: 2, 8, 7. (i) Which element is more reactive? (ii) Will A and B form ionic or covalent bonds when they combine? Explain using electron transfer or sharing. (iii) Predict the formula of the compound they would form. (i) B. It needs to gain only one electron to complete its octet, while A needs three, so B reacts more readily. (ii) Both have more than four valence electrons, so both tend to gain electrons; neither gives electrons away easily. They complete their octets by sharing, so the bonds are covalent. (iii) A needs three electrons and B needs one, so one A atom shares one electron with each of three B atoms. Formula: AB3. (A is phosphorus and B is chlorine; the compound is PCl3, phosphorus trichloride.)
5-Mark Questions
- Describe an activity with vinegar and baking soda to show that mass is conserved in a chemical reaction. Explain why an open set-up gives a different result from a closed one, and state the law with its proposer. Set-up 1: pour the baking soda into the open flask. There is brisk effervescence and the final reading is lower, because carbon dioxide escapes. Set-up 2: tie the balloon over the mouth of the flask, weigh, then lift the balloon so the baking soda falls in. Effervescence inflates the balloon and the final reading equals the initial one, because the gas is trapped and weighed. Conclusion: total mass of reactants equals total mass of products. Law of Conservation of Mass: matter can neither be created nor destroyed in a chemical reaction, proposed by Antoine Lavoisier in 1789.
- Differentiate between ionic and covalent compounds on the basis of (i) formation (ii) solubility in water and kerosene (iii) electrical conductivity in the solid state and in water (iv) melting and boiling points. Give two examples of each. Explain why an ionic compound conducts in water but not as a solid. (i) Ionic: transfer of electrons; covalent: sharing of electrons. (ii) Ionic: generally soluble in water, insoluble in kerosene and petrol; covalent: mostly insoluble in water, soluble in kerosene and petrol. (iii) Ionic: no conduction as a solid, conduction in water; covalent: no conduction in either (sugar solution does not conduct). (iv) Ionic: generally high; covalent: usually low. Examples: NaCl, copper sulfate (ionic); camphor, naphthalene (covalent). In the solid, ions are held in fixed positions by strong forces; in water they are free to move and carry charge.
Common Mistakes and Examiner Tips
- Saying the law “fails” when a gas escapes. A drop in mass in an open vessel means a gaseous product left the system. The law always holds; the measurement missed a product.
- Adding only some of the products. In a verification sum, add every reactant and every product listed, then write “mass of reactants = mass of products” as the final line. Missing the water in Example 9.1 breaks the balance.
- Scaling the wrong way in a ratio sum. For “46 g of sodium, ratio 23 : 35.5”, the factor is 46 ÷ 23 = 2, applied to 35.5. Dividing 46 by 35.5 is the usual slip.
- Drawing O2 with a single bond. Oxygen shares two electrons each, so O=O is a double bond. Nitrogen shares three each, a triple bond. Count the electrons each atom still needs.
- Using mono- for the first element. CO2 is carbon dioxide, never monocarbon dioxide. CO is carbon monoxide, never monooxide. H2S is hydrogen sulfide, never dihydrogen sulfide.
- Criss-crossing the signs. Only the numbers cross over, and charges are not written in the final formula. Write CaCl2, never Ca2+Cl2−.
- Forgetting to simplify. Mg2O2 must become MgO, and Ca2(CO3)2 must become CaCO3. An ionic formula shows the simplest ratio.
- Missing brackets on polyatomic ions. Write Al(OH)3, Mg(OH)2, Al2(SO4)3. Do not add brackets when there is only one polyatomic ion: NaOH, K2SO4.
- Mixing up ferrous and ferric, cuprous and cupric. -ous is the lower charge (Fe2+, Cu+) and -ic the higher (Fe3+, Cu2+). Ferric oxide is Fe2O3; cuprous oxide is Cu2O.
- Using “molecular mass” for an ionic compound, or missing the outer subscript. NaCl, Na2O and Ca(NO3)2 have formula unit masses. In Ca(NO3)2, multiply the whole nitrate mass (62 u) by 2 before adding calcium.
Quick Revision Points
- Law of Conservation of Mass: matter can neither be created nor destroyed in a chemical reaction; Antoine Lavoisier, 1789, the Father of Modern Chemistry.
- Activity 9.2: open set-up loses mass because CO2 escapes; with the balloon tied on, the mass stays the same.
- Law of Constant Proportions (Definite Proportions, Proust’s Law): elements in a compound combine in a fixed ratio by mass, whatever the source.
- Water: hydrogen : oxygen = 1 : 8 by mass; 9 g water gives 1 g H and 8 g O.
- Dalton, 1808: atoms are tiny, indivisible, not created or destroyed in reactions; same element, identical atoms; combine in simple whole number ratios; relative number and kinds of atoms are constant in a compound.
- Molecule: electrically neutral, more than one atom, independent existence, shows all properties of the substance. Helium exists as atoms.
- Chemical bond = force holding atoms together. Covalent = sharing; ionic = transfer.
- H-H, Cl-Cl, H-Cl: single bonds. O=O: double bond. N≡N: triple bond. Water: H-O-H.
- Prefixes: mono 1, di 2, tri 3, tetra 4, penta 5, hexa 6; no mono- on the first element; monoxide, pentoxide, tetroxide drop a vowel.
- Fewer than four valence electrons: usually lose them. More than four: usually gain or share.
- Na (2, 8, 1) → Na+ (11 p, 10 e); Cl (2, 8, 7) → Cl− (17 p, 18 e). Ionic bond = electrostatic attraction between them.
- Ionic compounds: generally soluble in water, insoluble in kerosene and petrol, conduct in water and when molten, not as solids, high melting points.
- Covalent compounds: mostly insoluble in water, soluble in kerosene and petrol, do not conduct, low melting points; sugar dissolves in water but gives no ions.
- Criss-cross: swap valency numbers as subscripts, divide by any common factor, bracket repeated polyatomic ions.
- Molecular mass: H2O 18 u, CO2 44 u, HNO3 63 u, CH4 16 u.
- Formula unit mass: Na2O 62 u, Ca(NO3)2 164 u, KCl 74.5 u, Mg(OH)2 58 u.
Weightage and Exam Pattern
Class 9 has no board exam. This chapter is assessed in school exams (periodic tests, half-yearly and annual papers), and the NCERT source gives no fixed mark allocation for it.
| Question type | What this chapter typically supplies |
|---|---|
| MCQ and one-word | Proposers of the laws, naming compounds with prefixes, charges of common ions, choosing the neutral combination of ions |
| Assertion-Reason | Conductivity of ionic compounds in solid and molten states, sugar solution, stability of noble gases, Dalton’s postulates |
| Short answer (2 to 3 marks) | Mass verification sums, fixed-ratio sums, formulae from pairs of ions, molecular mass and formula unit mass, forming a cation or an anion |
| Long answer (5 marks) | Describing Activity 9.2 or 9.3, formation of NaCl or H2O with explanation, comparing ionic and covalent compounds |
| Case-based | Open versus closed reactions, water from different sources, designing an ionic compound, Activity 9.4 observation tables |
This chapter is the foundation for Class 10 chemistry: valency, formula writing and bonding are used in every later chemistry chapter.
Class 9 Science · Chapter 9 – swipe through all 9 cards to understand the whole chapter.
Conservation of Mass
Matter can neither be created nor destroyed in a chemical reaction.
Proposed by Antoine Lavoisier in 1789.
- Open flask: CO2 escapes, so the balance reading falls.
- Balloon tied on: gas trapped, reading stays the same.
- 4.0 g + 2.92 g = 1.76 g + 0.72 g + 4.44 g = 6.92 g.
Constant Proportions
Elements in a compound combine in a fixed ratio by mass, whatever the source.
Also called the Law of Definite Proportions or Proust’s Law.
- 9 g of pure water always gives 1 g H and 8 g O.
- NaCl is 23 : 35.5, so 46 g Na needs 71 g Cl.
- Holds for compounds, not for mixtures.
Dalton’s Atomic Theory
Dalton’s postulates (1808) explain both laws using atoms.
A postulate is an assumption accepted as true without formal proof.
- Atoms cannot be created or destroyed in a reaction, only rearranged.
- Atoms of one element are identical in mass and chemical properties.
- Atoms combine in simple whole number ratios to form compounds.
Covalent Bonds
Atoms share electron pairs to reach an octet, or a duplet for hydrogen.
One shared pair is a single bond; two shared pairs a double bond.
- H2, Cl2 and HCl each share one pair of electrons.
- O2 shares two pairs because oxygen needs two electrons.
- Helium and neon are already stable and exist as atoms.
Naming Covalent Compounds
Prefixes show how many atoms of each element are present.
No mono- on the first element; second element ends in -ide.
- CO is carbon monoxide; N2O5 is dinitrogen pentoxide.
- SF6 is sulfur hexafluoride; PCl3 is phosphorus trichloride.
- H2S is hydrogen sulfide: hydrogen first takes no prefix.
Ionic Bonds
One atom loses electrons, another gains them, and the ions attract.
Ionic bond: electrostatic attraction between oppositely charged ions.
- Na⁺ has 11 protons and 10 electrons.
- In NaCl crystals each ion has six neighbours of opposite charge.
- Metals form cations; non-metals form anions.
Ionic vs Covalent Compounds
Solubility and conductivity tests separate the two kinds.
Ions must be free to move for a substance to conduct.
- Ionic: dissolve in water, conduct in solution and when molten.
- Covalent: camphor and naphthalene dissolve in kerosene, do not conduct.
- Sugar dissolves in water but gives no ions, so no current.
The Criss-Cross Method
Swap the charge numbers as subscripts, then simplify.
Total positive charge must equal total negative charge.
- Mg2O2 simplifies to MgO; Ca2(CO3)2 to CaCO3.
- Bracket a repeated polyatomic ion: Al(OH)3, Al2(SO4)3.
- Ferrous Fe2⁺, ferric Fe3⁺; cuprous Cu⁺, cupric Cu2⁺.
Molecular and Formula Unit Mass
Add the atomic masses of every atom in the formula.
Ionic compounds use formula unit mass, since they form no molecules.
- H2O = 18 u, HNO3 = 63 u, CH4 = 16 u.
- Na2O = 62 u, KCl = 74.5 u, Mg(OH)2 = 58 u.
- Ca(NO3)2 = 40 + 62 × 2 = 164 u.
📝 Practice Atomic Foundations of Matter - 10 questions
Concept-check questions · with answers & explanations
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Why B: Burning ethanol is a chemical reaction whose products are carbon dioxide and water vapour. In an open beaker both leave as gases, so nothing remains in the beaker. The mass of the escaped products, together with the oxygen used up from the air, still equals the mass of the ethanol and oxygen that reacted. If the burning were done in a closed container the reading before and after would be the same.
Why not A: Matter can neither be created nor destroyed in a chemical reaction; an empty beaker only shows that the products are not solids.
Why not C: Ethanol’s atoms are rearranged into carbon dioxide and water, and those atoms keep their mass; no substance becomes massless.
Why not D: The law applies to every chemical reaction; gases and liquids obey it as much as solids do.
Remember: No residue does not mean no mass; look for the gases that left.
Why A: The law says the elements in a compound combine in a fixed ratio by mass. Y’s ratio 8:2 divides by 2 to give 4:1, exactly X’s ratio. Both samples of the oxide have the same proportion of copper to oxygen by mass, which is what the law predicts, so the results justify it.
Why not B: The numbers differ but the ratio does not; 8:2 and 4:1 are the same proportion.
Why not C: The law applies to every compound, as Proust showed with copper carbonate and as cinnabar shows for mercury and sulfur.
Why not D: The law says the opposite; a given compound needs one particular ratio, and a different ratio would be a different compound or leftover reactant.
Remember: Simplify before you compare; same simplest ratio means the same compound.
Why A: The assertion is the chapter’s own example of the Law of Constant Proportions: water from rivers, borewells or the ocean, once purified, always has hydrogen and oxygen in the mass ratio 1:8. The reason quotes two of Dalton’s postulates correctly. If every water molecule has the same number and kinds of atoms (two hydrogen, one oxygen) and every atom of an element has the same mass, the mass ratio of hydrogen to oxygen must be the same in every sample. So R is true and it is exactly the explanation of A.
Why not B: R is not a side fact here; fixed numbers of fixed-mass atoms are precisely why the mass ratio is fixed.
Why not C: Both parts of R are postulates stated in the chapter, so R is true.
Why not D: The 1:8 mass ratio of water is the chapter’s standard result, true for water from any source.
Remember: Same atoms in the same numbers, each with a fixed mass, give the same mass ratio every time.
Why B: A has five valence electrons and needs three; B has seven and needs one. Both have more than four valence electrons, so neither donates; they share. Each B atom shares one pair with A, and A needs three such pairs, so three B atoms bond to one A atom, giving the covalent compound AB3. This mirrors phosphorus trichloride, PCl3.
Why not A: An ionic bond needs one atom to give electrons, which happens for atoms with fewer than four valence electrons; both A and B have more than four.
Why not C: A3B would have three A atoms around one B, but B needs only one electron and can share only one pair.
Why not D: The bond is covalent, and A3B has the counts reversed.
Remember: Both non-metals share; the atom needing three gets three partners that each need one.
Why C: The assertion is true: an ionic compound like copper sulfate conducts when molten because melting frees the ions, and does not conduct as a solid because the ions are held in fixed positions. The reason has the two states swapped: ions are fixed in the solid lattice and free in the molten state. So R is false.
Why not A: R contradicts the actual behaviour of the lattice, so it cannot explain A.
Why not B: R is not merely irrelevant, it is false.
Why not D: A is correct; molten ionic compounds conduct and solid ones do not.
Remember: Solid lattice locks ions; melting or dissolving frees them.
Why C: Two Fe3+ ions give 2 × 3 = 6+. Three O2- ions give 3 × 2 = 6-. The charges balance exactly, so the compound Fe2O3 (ferric oxide) is neutral and uses the stated totals.
Why not A: 2 Al3+ give 6+, but 3 Cl- give only 3-; six chloride ions would be needed, and the simplest formula is AlCl3.
Why not B: 3 Mg2+ give 6+, but one PO4 3- gives only 3-; two phosphate ions are needed, as in Mg3(PO4)2.
Why not D: 3 Ca2+ give 6+, but 2 SO4 2- give only 4-; the balanced formula is CaSO4 with one of each.
Remember: Multiply each ion’s charge by its count; the two totals must be equal and opposite.
Why B: Vinegar and baking soda (sodium hydrogencarbonate) react to give carbon dioxide and other substances. In the open flask the carbon dioxide escapes, so the gas is no longer on the balance and the reading drops. When the balloon is tied to the mouth of the flask (set-up 2) the gas is trapped and the initial and final readings match, showing that mass was conserved all along.
Why not A: Matter can neither be created nor destroyed in a chemical reaction; the missing mass left as gas, it was not destroyed.
Why not C: The effervescence is the gas being released; there is no significant loss of vinegar by evaporation in a brief reaction, and the drop happens at once.
Why not D: A digital balance has an uncertainty of only ±1 in the last digit; it does not change its sensitivity because of a reaction.
Remember: Open flask plus a gas product means a lower reading; close the system and the mass is conserved.
Why C: Step 1 (Law of Constant Proportions): sulfur to oxygen is 40:60, that is 2:3 by mass in every sample. Step 2 (proportion): 40 g of sulfur pairs with 60 g of oxygen, so 20 g of sulfur pairs with (60 ÷ 40) × 20 = 30 g of oxygen.
Why not B: 20 g would make the ratio 1:1, but the compound is 40:60, so oxygen must exceed sulfur.
Why not A: 12 g is 60 % of 20 g, treating the sulfur mass as the whole sample; the 60 % refers to the whole compound, not to the sulfur alone.
Why not D: 50 g would make the sample 20 g sulfur in 70 g, about 29 % sulfur, not 40 %.
Remember: Convert the percentages to a ratio, then scale that ratio to the given mass.
Why D: A compound always has the same relative number and kinds of atoms, and every atom of an element has the same mass. Fixed numbers of fixed-mass atoms give a fixed mass ratio of elements, whatever the source of the compound. This is how the atomic theory explains why substances combine in fixed proportions.
Why not B: Indestructible atoms explain conservation of mass, not the fixed composition of a compound.
Why not C: Saying atoms take part in reactions does not fix how many of each kind are in a compound.
Why not A: That matter is made of atoms is the general starting point; it does not by itself fix the ratio within a compound.
Remember: Fixed number and kind of atoms in a compound means a fixed ratio by mass.
Why D: Oxygen has six valence electrons and needs two more for an octet. A hydrogen atom has one electron and needs one more for its duplet, so it can share only one pair. Oxygen therefore shares one pair with each of two hydrogen atoms, and the molecule is H2O with two single covalent bonds.
Why not B: Hydrogen appears singly in hydrogen chloride, HCl; the count depends on what the partner atom needs.
Why not C: Water is a covalent compound formed by sharing; no electrons are transferred, and oxygen gains rather than gives electrons in any case.
Why not A: Mass has nothing to do with it; the number of atoms is fixed by the electrons each atom needs.
Remember: Count the electrons each atom needs; the shortfall decides how many partners it takes.
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Frequently Asked Questions
It states that matter can neither be created nor destroyed in a chemical reaction, so the total mass of reactants equals the total mass of products. Antoine Lavoisier proposed it in 1789. If mass seems to drop in an open vessel, a gas has escaped; in a closed set-up the balance reading stays the same.
The reaction produces carbon dioxide gas. In the open flask the gas escapes into the air, so its mass is not weighed and the final reading is lower. When the balloon of baking soda is tied over the mouth of the flask, the gas is trapped, the balloon inflates, and the final reading equals the initial reading.
Proust’s law says that in any compound the elements combine in a fixed ratio by mass, whatever the source. Purified water from a river, a borewell or the ocean always has hydrogen and oxygen in the mass ratio 1 : 8, so 9 g of water gives 1 g of hydrogen and 8 g of oxygen. The law holds for compounds, not mixtures.
Dalton said atoms cannot be created or destroyed in a reaction and are only rearranged, so the total mass stays the same. He also said atoms of an element have a fixed mass and combine in simple whole number ratios, with the relative number and kinds of atoms constant in a compound, so the mass ratio in a compound is fixed.
A covalent bond forms when atoms share pairs of electrons, as in H₂, O₂, HCl and H₂O. An ionic bond forms when one atom transfers electrons to another, making a cation and an anion that are held by electrostatic attraction, as in NaCl. Atoms with fewer than four valence electrons usually lose them; atoms with more than four usually gain or share.
Salt is an ionic compound. When it dissolves, its Na⁺ and Cl⁻ ions become free to move and carry the current. Sugar is a covalent compound that dissolves as neutral molecules and gives no ions, so its solution does not conduct. Solid salt also does not conduct, because its ions are held in fixed positions.
Write the cation symbol first and the anion second, put the charge numbers below them, and cross the numbers over as subscripts. Divide by any common factor and put brackets around a polyatomic ion that appears more than once. Al³⁺ and O²⁻ give Al₂O₃; Mg²⁺ and O²⁻ give Mg₂O₂, written MgO; Al³⁺ and OH⁻ give Al(OH)₃.
Molecular mass is used for covalent compounds and is the sum of the atomic masses of all atoms in one molecule, such as H₂O = 18 u. Ionic compounds form crystals and have no separate molecules, so we use formula unit mass, the sum of atomic masses in the simplest ratio of ions, such as Na₂O = 62 u and Ca(NO₃)₂ = 164 u.
