NCERT Solutions Class 9 Science Exploration Chapter 09 Atomic Foundations of Matter

Question 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?
Answer: Element A has one electron in its outermost shell and will lose that single electron to reach stability. This occurs because removing one electron is simpler than gaining seven additional electrons. When A loses this electron, it becomes a positively charged particle (cation) with a +1 charge, represented as A⁺. Element B possesses six electrons in its second shell. The second shell can accommodate a maximum of eight electrons, so B requires two more electrons to achieve stability. Since B accepts two electrons, it transforms into a negatively charged particle (anion) with a -2 charge, written as B²⁻. When A and B react together, A contributes one electron while B needs two electrons, meaning two atoms of A bond with one atom of B. An Ionic Bond forms through this electron transfer process from A to B. Applying the criss-cross method reveals that A has a valency of 1 while B has a valency of 2, giving the compound formula A₂B.
In simple words: Element A loses one electron to become A⁺. Element B takes two electrons to become B²⁻. Two A atoms join with one B atom to make A₂B using an ionic bond.

Exam Tip: Always count valence electrons first to determine whether an element will donate or accept electrons, then use the criss-cross method to find the correct formula.

Question 2. An element X has six electrons in its outer shell and forms a diatomic molecule. (i) Why would that be so? (ii) What kind of bond would it form? (iii) Draw the structure of the molecule it would form. (iv) A certain other element Y has two electrons in its second shell. Draw the structure of the molecule that X would form with Y.
Answer: Element X has six valence electrons and needs two additional electrons to complete its octet. Since X is a non-metal, it achieves stability through electron sharing with another atom of the same element, forming a diatomic molecule (X₂). Each atom shares two electrons with the other atom, creating a Double Covalent Bond, similar to the structure found in oxygen (O₂). The structure can be represented as X=X, showing two shared electron pairs between the atoms. Element Y has two electrons in its second shell, meaning its valency is 2 (it would need six more to fill the octet, but sharing two is more straightforward). When X and Y combine, Y shares two electrons with X - just like the arrangement in water (H₂O). The formula formed would be YX. If Y behaves like a metal similar to magnesium, an ionic bond forms; if Y resembles beryllium-type elements, a covalent bond develops. If Y itself is a non-metal like oxygen, the compound formed would be XY with a double bond.
In simple words: X needs two more electrons, so two X atoms share two electron pairs, making a double bond. X and Y can form different types of bonds depending on what Y is.

Exam Tip: Recognize that non-metals form covalent bonds by sharing electrons, and the type of bond depends on how many electron pairs are shared.

Question 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) 2 Fe³⁺ and 3 O²⁻
In simple words: You need to check each option to see if positive and negative charges match. For option (iii): 2 Fe³⁺ gives 2 × 3 = 6 positive charge, and 3 O²⁻ gives 3 × 2 = 6 negative charge. They balance perfectly.
To verify all options: (i) gives 6+ and 3-, which do not match. (ii) gives 6+ and 3-, which do not match. (iv) gives 6+ and 4-, which do not match.

Exam Tip: Always multiply the ion charge by its quantity to find total positive and negative charges, then confirm they are equal for a balanced ionic compound.

Question 4. 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.
Answer: Statement (i) is incorrect. The right statement is: Elements are made up of atoms, while compounds are made up of molecules (which contain atoms of different elements). Statement (ii) is incorrect. The correction is: A molecule of a compound is always made up of two or more atoms of different kinds of elements, not the same kind. Statement (iii) is incorrect. The correction is: One molecule of nitrogen gas (N₂) contains two nitrogen atoms, not three. Statement (iv) is correct as written. Water (H₂O) consists of two hydrogen atoms covalently bonded with one oxygen atom, and this statement accurately describes this arrangement.
In simple words: Atoms make elements. Molecules (with different atoms) make compounds. A nitrogen gas molecule has two nitrogen atoms, not three. Water's description is correct.

Exam Tip: Remember the distinction between atoms and molecules: atoms are the basic units of elements, while molecules are combinations that form compounds or element molecules.

Question 5. Write the chemical formulae for the following compounds. (i) Aluminium nitrate (ii) Calcium oxide (iii) Ferric oxide
Answer: Aluminium nitrate is represented as Al(NO₃)₃. Calcium oxide is written as CaO. Ferric oxide has the formula Fe₂O₃.
In simple words: Use the criss-cross method with ion charges to find each formula. Al³⁺ with NO₃⁻ gives Al(NO₃)₃. Ca²⁺ with O²⁻ gives CaO. Fe³⁺ with O²⁻ gives Fe₂O₃.

Exam Tip: Always identify the charge on each ion first, then apply the criss-cross method. Use brackets for polyatomic ions when the subscript is greater than 1.

Question 6. 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⁻
Answer: Ca²⁺ and Br⁻ combine to produce CaBr₂ (Calcium bromide). Al³⁺ and CO₃²⁻ produce Al₂(CO₃)₃ (Aluminium carbonate). K⁺ and SO₄²⁻ produce K₂SO₄ (Potassium sulfate). NH₄⁺ and Cl⁻ produce NH₄Cl (Ammonium chloride).
In simple words: Match cation charge with anion charge using the criss-cross method. Ca²⁺ needs 2 Br⁻ ions. Al³⁺ needs 3 CO₃²⁻ ions. K⁺ needs 2 SO₄²⁻ ions. NH₄⁺ needs 1 Cl⁻ ion.

Exam Tip: Write the cation first, then the anion. Cross the numbers as subscripts, reduce by any common factor, and use brackets around polyatomic ions when needed.

Question 7. Which of the following, in Fig. 9.18, correctly represents Cl⁻ ion (Atomic number of chlorine = 17).
Answer: Chlorine has an atomic number of 17, giving it an electronic configuration of 2, 8, 7 across shells K, L, and M. When chlorine accepts one electron to form the Cl⁻ ion, the M shell achieves completion with 8 electrons. The resulting configuration of Cl⁻ is 2, 8, 8, with a total of 18 electrons. Diagram (iii) correctly shows this electron arrangement for the Cl⁻ ion.
In simple words: Chlorine gains one electron to fill its outer shell to 8 electrons. This gives Cl⁻ a total of 18 electrons arranged as 2, 8, 8.

Exam Tip: Count the total electrons in the ion (atomic number plus gained electrons or minus lost electrons) and distribute them into K, L, M shells respecting the rule that each shell holds a maximum number of electrons.

Question 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.
Answer: For ammonium nitrate (NH₄NO₃): Adding the atomic masses - nitrogen contributes 14 u, hydrogen contributes 1 u each, and oxygen contributes 16 u each. Calculation: (14 × 1) + (1 × 4) + (14 × 1) + (16 × 3) = 14 + 4 + 14 + 48 = 80 u. For phosphoric acid (H₃PO₄): With hydrogen at 1 u, phosphorus at 31 u, and oxygen at 16 u each. Calculation: (1 × 3) + (31 × 1) + (16 × 4) = 3 + 31 + 64 = 98 u. For sodium hydrogencarbonate (NaHCO₃): With sodium at 23 u, hydrogen at 1 u, carbon at 12 u, and oxygen at 16 u each. Calculation: 23 + 1 + 12 + (16 × 3) = 23 + 1 + 12 + 48 = 84 u.
In simple words: Multiply the atomic mass of each element by how many atoms are in the formula, then add all the results together.

Exam Tip: Be careful to count every atom in the formula, especially those within parentheses like in NH₄NO₃. Use the atomic mass values provided in your periodic table.

Question 9. Write the formulae for the compounds formed by the reaction of: (i) Magnesium and nitrogen (ii) Lithium and nitrogen (iii) Sodium and sulfur (iv) Aluminium and oxygen
Answer: Magnesium (Mg²⁺) reacts with nitrogen (N³⁻) using the criss-cross method to give Mg₃N₂ (Magnesium nitride). Lithium (Li⁺) reacts with nitrogen (N³⁻) to produce Li₃N (Lithium nitride). Sodium (Na⁺) reacts with sulfur (S²⁻) to form Na₂S (Sodium sulfide). Aluminium (Al³⁺) reacts with oxygen (O²⁻) to create Al₂O₃ (Aluminium oxide).
In simple words: Identify the charge on each element, then apply criss-cross to find how many atoms of each combine to balance charges.

Exam Tip: Remember that metal atoms form cations (lose electrons) and non-metal atoms form anions (gain electrons). The criss-cross method swaps the charge numbers as subscripts.

Question 10. Complete the Table 9.3 by writing the formulae of the compounds formed by the cations on the left and the anions at the top. LiNO₃ is given as an example.
Answer:

In simple words: For each box, match the cation (left) with the anion (top) using the criss-cross method. The charge numbers swap to become subscripts.

Exam Tip: Always apply the criss-cross rule systematically. Remember to use brackets around polyatomic ions (like NO₃, SO₄, PO₄) when the subscript is 2 or higher.

Question 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.
Answer: The total mass of the starting materials is found by adding 5.3 g plus 6.0 g, which equals 11.3 g. The total mass of the products is calculated by adding 2.2 g plus 0.9 g plus 8.2 g, which also equals 11.3 g. Since the mass before the reaction equals the mass after the reaction (both 11.3 g), the Law of Conservation of Mass is confirmed. In this chemical reaction, matter is neither created nor destroyed - it simply changes form.
In simple words: Add up all starting materials: 5.3 + 6.0 = 11.3 g. Add up all products: 2.2 + 0.9 + 8.2 = 11.3 g. They match, so the law is proven correct.

Exam Tip: Always arrange your calculation clearly with total reactants on one side and total products on the other. Show that they are equal to verify the law of conservation of mass.

Question 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.
Answer: The atomic number equals the number of protons, which is 11. The mass number is determined by adding protons and neutrons: 11 + 12 = 23. Since the species has 11 protons but only 10 electrons, there is an imbalance. The positive charge from 11 protons exceeds the negative charge from 10 electrons. The net charge is 11 minus 10, which gives +1. Therefore, this species is a Cation with a charge of +1. To find the electronic configuration, count the total electrons (10) and distribute them into shells. The K shell holds 2 electrons and the L shell holds 8 electrons, giving the configuration 2, 8. Sodium (Na) has an atomic number of 11. This species with a charge of +1 is the Sodium ion, written as Na⁺.
In simple words: Atomic number = protons = 11. Mass number = protons + neutrons = 23. More protons than electrons means it is a cation with +1 charge. It is the sodium ion, Na⁺.

Exam Tip: Remember that atomic number always equals the number of protons. The mass number is the sum of protons and neutrons. Count electrons to determine if an ion is positive (cation) or negative (anion).

Question 13. Two elements, A and B, have the following configurations - A: 2, 8, 5 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.
Answer: Element A possesses 5 valence electrons and requires 3 additional electrons to complete its octet. Element B possesses 7 valence electrons and requires only 1 additional electron to achieve stability. Since B requires fewer electrons and has a stronger pull on electrons, B is more reactive (similar to chlorine, which is very reactive). When considering bonding between A and B: A has 5 valence electrons and B has 7, meaning neither has an extremely low electron count (less than 4) to simply transfer electrons. However, B needs just 1 electron while A can share electrons. If A shares with B - since A needs 3 and B needs 1 - A would share with 3 atoms of B to achieve full stability. This creates a Covalent Bond (electron sharing between non-metals). Note that A has 5 valence electrons like nitrogen and P, while B has 7 like chlorine and fluorine. Both are non-metals, confirming covalent bonding occurs through sharing. Using the criss-cross method: A requires 3 electrons (valency = 3) and B requires 1 electron (valency = 1), producing the formula AB₃ (similar to NCl₃, nitrogen trichloride).
In simple words: B needs fewer electrons so it is more reactive. A and B are both non-metals so they share electrons, making a covalent bond. A needs 3, B needs 1, so the formula is AB₃.

Exam Tip: Elements needing fewer electrons to fill their outer shell are generally more reactive. Non-metals sharing electrons form covalent bonds, while metals and non-metals form ionic bonds through electron transfer.

Question 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.
Answer: (iii) A is true, but R is false.
In simple words: Copper sulfate does conduct electricity when melted, which makes statement A correct. However, statement R gets the explanation backwards - ions are actually fixed in the solid state, not the molten state. In the molten state, ions move freely.

Exam Tip: Be careful with reason-based questions. The assertion can be true while the explanation is incorrect. Always verify both the statement AND the reasoning separately.

Question 15. The species ²⁷Al, ⁸⁰Br⁻ and ²⁰¹Hg²⁺ have 13, 35 and 80 protons, respectively. How many electrons and neutrons do they have?
Answer:

In simple words: For a neutral atom, electrons equal protons. For an anion (negative), add the charge number to protons. For a cation (positive), subtract the charge number from protons. Neutrons = mass number minus protons.

Exam Tip: Use the mass number (shown as a superscript) and atomic number (protons) to find neutrons. Adjust electrons based on the ion's charge (+ means fewer electrons, - means more electrons).

Extra Questions for Exam Preparation

Very Short Answer Type Questions

Question 1. State the Law of Conservation of Mass. Who proposed it and when?
Answer: In any chemical process, matter cannot be generated or destroyed; it only transforms. The total mass of reactants invariably equals the total mass of products. Antoine Lavoisier first put forward this law in 1789 and earned recognition as the founder of contemporary chemistry.
In simple words: In a chemical reaction, you do not lose or gain matter - the starting materials weigh the same as the ending materials.

Exam Tip: Always mention that this law applies only in closed systems where no material escapes or enters.

Question 2. State the Law of Constant Proportions.
Answer: In any compound formed through the combination of two or more elements, those elements unite in a consistent mass ratio that does not vary based on where the compound originates or how it is prepared. This principle is also known as Proust's Law.
In simple words: The elements in a compound always mix in the same amounts no matter how you make the compound or where you find it.

Exam Tip: This law applies to compounds only, not to mixtures - compounds have fixed proportions, but mixtures do not.

Question 3. Give two postulates of Dalton's Atomic Theory.
Answer: One fundamental idea in Dalton's theory states that all substances consist of very small particles known as atoms, which take part in chemical transformations. Another core concept is that atoms are indivisible units - they cannot be separated, generated, or removed during a chemical reaction. Dalton put forth six postulates in 1808 that formed the foundation for explaining the two basic laws of chemistry.
In simple words: All matter is made of atoms. Atoms cannot be broken apart, made, or destroyed in chemical reactions.

Exam Tip: Know all six postulates of Dalton's theory, not just two - examiners often ask for multiple postulates.

Question 4. Define a molecule.
Answer: A molecule represents a neutral particle made of multiple atoms that can remain by itself and displays the complete characteristics of that material. Examples include H₂ (the hydrogen molecule) and H₂O (the water molecule).
In simple words: A molecule is two or more atoms stuck together that can exist on its own and shows the properties of a substance.

Exam Tip: Note that some elements exist as single atoms (noble gases) and are not called molecules - only entities with 2+ atoms joined together are molecules.

Question 5. What is a covalent bond? Give one example.
Answer: A covalent bond develops through the sharing of one or multiple electron pairs among atoms to create steady electron setups. As an illustration, in the H₂ molecule, two hydrogen atoms each share one electron to produce a single covalent bond denoted by H-H.
In simple words: A covalent bond is when atoms share electrons with each other to become stable.

Exam Tip: Covalent bonds occur between non-metals. Always be ready to show the electron sharing using Lewis dot structures if asked.

Question 6. What is an ionic bond? Give one example.
Answer: An ionic bond represents the electrostatic attraction that develops between ions carrying opposite charges (a positively charged cation and a negatively charged anion) as a result of electron movement. As an instance, in NaCl, sodium loses a single electron to chlorine, producing Na⁺ and Cl⁻, which remain bound together through the ionic bond in the compound.
In simple words: An ionic bond forms when one atom gives electrons to another atom, and the opposite charges pull them together.

Exam Tip: Ionic bonds happen between metals and non-metals. The metal loses electrons to become a cation, and the non-metal gains them to become an anion.

Question 7. Distinguish between a cation and an anion.
Answer: A cation is an ion carrying a positive charge, formed when an atom releases one or multiple electrons. An anion is an ion carrying a negative charge, formed when an atom receives one or multiple electrons. Illustrations include Na⁺ functioning as a cation and Cl⁻ functioning as an anion. Both kinds are referred to collectively as ions.
In simple words: A cation is positive (lost electrons), and an anion is negative (gained electrons).

Exam Tip: Cations are typically metals, and anions are typically non-metals. The suffix "-on" helps remember: cation (ca-tion) = positive, anion (an-ion) = negative.

Question 8. Why does ionic compound sodium chloride not conduct electricity in the solid state?
Answer: In solid NaCl, the Na⁺ and Cl⁻ ions occupy unchanging spots inside the crystal framework, held in place by intense electrostatic attractions. Because the ions cannot roam freely throughout the solid arrangement, electric current cannot flow. Movement of ions is restricted to when the compound dissolves in water, which then permits conductivity to occur.
In simple words: Solid NaCl does not conduct electricity because the ions are stuck in place in the crystal. They need to be free to move to carry an electric current.

Exam Tip: Remember that electrical conductivity requires moving charged particles. When ions are locked in a fixed crystal structure, they cannot move, so conduction is impossible.

Question 9. What is the difference between molecular mass and formula unit mass?
Answer: Molecular mass represents the sum of atomic masses from all atoms in one molecule - this measure pertains to covalent compounds. Formula unit mass represents the sum of atomic masses from all atoms in a formula unit - this measure applies to ionic compounds, which arrange into crystal grids instead of separate molecules.
In simple words: Molecular mass is for compounds with molecules (covalent). Formula unit mass is for compounds with crystal patterns (ionic).

Exam Tip: The calculation method is the same for both, but the terminology differs based on the compound type.

Question 10. Name the covalent compound with formula SF₆ using IUPAC prefix system.
Answer: SF₆ receives the name sulfur hexafluoride. The word "hexa" points to the presence of 6 fluorine atoms. The second element (fluorine) transforms to end with -ide, turning into fluoride, and when combined with "hexa" yields hexafluoride. The initial element (sulfur) retains its original name without receiving a prefix.
In simple words: Count the atoms of the second element and use a prefix. The second element gets -ide added. First element stays as is.

Exam Tip: Learn the prefixes: mono(1), di(2), tri(3), tetra(4), penta(5), hexa(6), hepta(7), octa(8). Remember that "a" or "o" is often dropped before another vowel.

Question 11. Write the chemical formula of calcium carbonate using the criss-cross method.
Answer: The calcium ion carries a charge of 2 (Ca²⁺); the carbonate ion carries a charge of 2 (CO₃²⁻). By crossing the charges, we would get Ca₂(CO₃)₂. Dividing both subscripts by the common factor 2 yields the most basic formula: CaCO₃. The charges do not appear in the completed formula.
In simple words: Ca²⁺ with CO₃²⁻ would give Ca₂(CO₃)₂, but since both numbers divide by 2, the final answer is CaCO₃.

Exam Tip: Always reduce your formula to the simplest whole number ratio. If both subscripts share a common factor, divide them both.

Question 12. Why does sugar dissolve in water but not conduct electricity?
Answer: Sugar functions as a covalent material. When introduced to water, it does not break into separate charged particles - it persists as sugar particles mixed throughout the solution. Electric current flow demands freely roaming, charged particles (ions), and since no ions form when sugar dissolves, the sugar liquid fails to enable conductivity.
In simple words: Sugar dissolves in water but stays as sugar molecules, not ions. Electricity needs moving ions to flow, so sugar solution does not conduct.

Exam Tip: Solubility and electrical conductivity are different properties - just because something dissolves does not mean it conducts electricity.

Question 13. What type of bond is present in an oxygen molecule (O₂)? How many pairs of electrons are shared?
Answer: The oxygen molecule (O₂) features a double covalent bond. Since oxygen carries 6 valence electrons and requires 2 additional ones, two oxygen atoms each share 2 electrons, generating two shared electron pairs. The structure shows the double bond as O=O, represented with two lines connecting the atoms.
In simple words: Oxygen atoms each share two electrons with each other, making two shared pairs - this is a double bond shown as O=O.

Exam Tip: Remember: single bond = 1 shared pair, double bond = 2 shared pairs, triple bond = 3 shared pairs.

Question 14. Write the formula of magnesium hydroxide. Why are brackets used?
Answer: Magnesium hydroxide carries the formula Mg(OH)₂. The criss-cross procedure applied to Mg²⁺ (with charge 2) and OH⁻ (with charge 1) produces Mg(OH)₂ - indicating that each magnesium ion pairs with two hydroxide ions. Brackets prove essential around the polyatomic ion (OH) to show that the subscript 2 relates to the complete OH entity, rather than being applied to just the oxygen.
In simple words: Brackets show that the subscript 2 applies to the whole OH group, not just one atom in it.

Exam Tip: Whenever a polyatomic ion has a subscript greater than 1, it must be enclosed in brackets to avoid confusion about which atom the subscript modifies.

Question 15. Calculate the molecular mass of water (H₂O).
Answer: The atomic mass measurements are H = 1 u and O = 16 u. The molecular mass formula for H₂O calculates as (1 u × 2) + (16 u × 1) = 2 + 16 = 18 u. Water consists of 2 hydrogen atoms bonded to 1 oxygen atom. The molecular mass (18 u) characterizes covalent materials whose atoms form separate molecules.
In simple words: Water has two hydrogen atoms (2 × 1 = 2) and one oxygen atom (1 × 16 = 16), so 2 + 16 = 18 u total.

Exam Tip: Molecular mass applies to covalent compounds. Always check the formula carefully to count how many atoms of each element are present.

Short Answer Type Questions

Question 1. Explain why the apparent mass decreases in Experimental Set-up 1 of Activity 9.2 but not in Set-up 2. Does this violate the Law of Conservation of Mass?
Answer: In Set-up 1, the balloon stands apart from the container - as baking soda interacts with vinegar, CO₂ gas forms and spreads into the surrounding environment. Because the released gas is not weighed during the measurement, the final result appears to be less than the starting amount. In Set-up 2, a balloon remains fastened to the flask's opening. The CO₂ that forms fills the balloon but stays contained within the closed arrangement. The overall mass (including the flask, balloon, and all substances created) stays equal to the mass at the beginning. This situation does NOT break the Law of Conservation of Mass. In each scenario, the actual combined mass remains intact. Set-up 1 simply neglects to account for the gas that escaped. The rule continues to be valid whenever the arrangement is sealed and keeps track of all outcomes.
In simple words: In Set-up 1, gas escapes and is not weighed, making it appear mass is lost. In Set-up 2, the gas stays in the balloon, so total mass does not change. The law is not broken - we just did not weigh everything in Set-up 1.

Exam Tip: The law of conservation of mass always holds true - what appears to break the law is usually an open system where a gas escapes unweighed.

Question 2. Why does the Law of Constant Proportions apply to compounds but not to mixtures?
Answer: In compounds, elements link together chemically in a particular, unchanging proportion determined by their bonding behaviors and atomic nature. This ratio stays the same no matter what the source is or what method was employed to create the compound. Take water as an instance - it invariably holds hydrogen and oxygen in a 1:8 mass proportion. In mixtures, components combine only in a physical manner without establishing chemical connections. They can be combined in any quantity - for instance, air can hold changing levels of nitrogen and oxygen, and salt can break down in water at different strengths. No fixed proportion is necessary for physical mixing. Consequently, the Law of Constant Proportions represents a characteristic of materials created through chemical combination (compounds) and not through physical blending (mixtures). This principle is why the regulation is sometimes referred to as the Law of Definite Proportions - the percentages in a compound remain consistent and do not shift.
In simple words: Compounds have fixed ratios because atoms bond in definite ways. Mixtures can have any ratio because they are just physically mixed, not chemically bonded.

Exam Tip: Always distinguish between compounds (fixed chemical ratio) and mixtures (variable physical proportions) when answering this question.

Question 3. Explain the formation of a water molecule (H₂O) with the help of electron sharing. What type of bond is formed?
Answer: Oxygen has atomic number 8 with electron arrangement 2, 6 - holding 6 valence electrons yet requiring 2 more to finish its octet. Hydrogen has atomic number 1 with one electron in the K-shell needing one additional to finish its duplet. To fulfill both requirements, a pair of hydrogen atoms each transfer one electron to the oxygen atom. The oxygen atom gives back one electron to each hydrogen atom. The outcome produces two single covalent bonds within the H₂O water molecule. Written as H-O-H, each dash shows a shared electron pair (single covalent bond). The molecule holds no electric charge and behaves stably. Since the atoms reach stability through sharing, water functions as a covalent material. Both connections in H₂O work as single covalent bonds created by sharing of electron pairs.
In simple words: Hydrogen needs one more electron, and oxygen needs two more. Two hydrogen atoms each share one electron with oxygen, forming H-O-H. This makes water a covalent compound.

Exam Tip: Be able to explain electron sharing using Lewis dot structures and electron configuration concepts.

Question 4. Explain the formation of sodium chloride (NaCl) through ionic bonding. Why is NaCl electrically neutral overall?
Answer: Sodium has atomic number 11 with electron setup 2, 8, 1 - one valence electron that it readily releases to reach a balanced octet in the L-shell. After sodium transfers this electron, it transforms into Na⁺ cation (11 protons, 10 electrons, net charge +1). Chlorine carries atomic number 17 with setup 2, 8, 7 - holding seven valence electrons and needing one additional to finish its octet. Chlorine takes the electron given off by sodium, converting to Cl⁻ anion (17 protons, 18 electrons, net charge -1). Na⁺ and Cl⁻ attract through intense electrostatic interaction - this is the ionic bond. They combine to create NaCl (sodium chloride). NaCl shows electrical neutrality in total since the +1 charge of Na⁺ precisely cancels out the -1 charge of Cl⁻. Overall positive charge = overall negative charge = neutral compound result.
In simple words: Sodium loses one electron and becomes +1. Chlorine gains that electron and becomes -1. The +1 and -1 cancel out, making NaCl neutral.

Exam Tip: When explaining ionic bonding, clearly show electron transfer and how opposite charges balance to create a neutral compound.

Question 5. Compare the properties of ionic and covalent compounds under three headings: solubility, electrical conductivity, and melting/boiling points.
Answer: Solubility: Ionic substances like NaCl and CuSO₄ typically break down easily in polar liquids such as water but do not break down in non-polar organic liquids like kerosene and petrol. Covalent compounds like camphor and naphthalene typically do not break down in water yet do break down in non-polar organic liquids. Electrical Conductivity: Ionic compounds fail to conduct electricity when solid (ions remain attached to the lattice) but conduct when spread in water (ions move freely). Covalent compounds do not conduct electricity even when dissolved since they do not break down into ions. One exception is sugar - it dissolves in water but still does not conduct. Melting and Boiling Points: Ionic compounds display elevated melting and boiling points resulting from intense electrostatic forces holding oppositely charged ions. Covalent compounds typically show reduced melting and boiling points due to much less strong intermolecular interactions as opposed to ionic connections.
In simple words: Ionic compounds dissolve in water and conduct when dissolved. Covalent compounds do not dissolve easily in water or conduct electricity. Ionic compounds have high melting points, covalent ones have low ones.

Exam Tip: Create a comparison table to memorize these properties clearly - it helps during exams to have a visual summary.

Question 6. Write the chemical formulae of the following ionic compounds using criss-cross method: (a) Aluminium oxide, (b) Calcium carbonate, (c) Magnesium hydroxide.
Answer: For aluminium oxide: Al³⁺ holds a charge of 3 and O²⁻ holds a charge of 2. Criss-crossing 3 and 2 provides Al₂O₃. The formula is Al₂O₃. For calcium carbonate: Ca²⁺ holds a charge of 2 and CO₃²⁻ holds a charge of 2. Equal charges criss-cross to provide Ca₂(CO₃)₂, then divide by 2 to get CaCO₃. The formula is CaCO₃. For magnesium hydroxide: Mg²⁺ carries a charge of 2 and OH⁻ carries a charge of 1. Criss-crossing produces Mg(OH)₂. Brackets are essential since subscript 2 relates to the complete polyatomic OH group. The formula is Mg(OH)₂. Note: In all formulas, charges on ions are not written in the final formula.
In simple words: Cross the numbers as subscripts. Reduce by common factors. Use brackets around polyatomic ions when needed.

Exam Tip: Practice the criss-cross method repeatedly until it becomes automatic - it is fundamental to writing correct ionic formulas.

Question 7. How does Dalton's Atomic Theory explain both the Law of Conservation of Mass and the Law of Constant Proportions?
Answer: Dalton's viewpoint held that atoms represent basic particles incapable of being broken down, generated, or removed in chemical processes - they simply rearrange. This immediately accounts for the Law of Conservation of Mass: because no atoms are generated or removed, the complete mass of all atoms prior to reacting remains equal to the complete mass after reacting. Regarding the Law of Constant Proportions, Dalton's perspective included the idea that atoms making up a particular element remain the same in their mass and that combinations form when atoms link in constant whole number proportions. Since atoms forever link in the same fixed proportion, the mass proportion of components in any compound stays invariable, no matter the origin or preparation approach. Accordingly, Dalton's framework supplied the particle-level description for both experimental rules: indivisibility of atoms elucidates conservation of mass, and fixed whole-number combining proportions explain constant proportions.
In simple words: Atoms do not change, so mass is conserved. Atoms always combine in the same ratios, so proportions are constant.

Exam Tip: Connect Dalton's postulates directly to each law - show how each postulate supports the law being explained.

Question 8. Name the following covalent compounds using IUPAC prefix rules: (a) CO, (b) CO₂, (c) N₂O₄, (d) PCl₃, (e) SF₆.
Answer: CO receives the name carbon monoxide. The word "mono" shows 1 oxygen atom. The prefix mono- applies to the second element even when left out for the initial one. CO₂ receives the name carbon dioxide. The word "di" shows 2 oxygen atoms. The prefix "di" keeps its "i" before oxide. N₂O₄ receives the name dinitrogen tetroxide. "Di" indicates 2 nitrogen; "tetra" indicates 4 oxygen, yet the ending "a" of "tetra" gets removed prior to "oxide" - producing tetroxide rather than tetraoxide. PCl₃ receives the name phosphorus trichloride. "Tri" shows 3 chlorine atoms. The element chlorine turns to chloride for the second element. SF₆ receives the name sulfur hexafluoride. "Hexa" shows 6 fluorine atoms. The ending "a" of "hexa" stays intact because fluoride begins with a consonant.
In simple words: Use the right prefix for each number, and remember vowel-dropping rules before a vowel.

Exam Tip: Memorize the prefixes and vowel-dropping rules well - they are essential for naming covalent compounds correctly.

Question 9. Calculate the molecular mass of nitric acid (HNO₃) and the formula unit mass of calcium nitrate Ca(NO₃)₂.
Answer: For nitric acid (HNO₃): Atomic masses are H = 1 u, N = 14 u, O = 16 u. Molecular mass calculation: (1 × 1) + (14 × 1) + (16 × 3) = 1 + 14 + 48 = 63 u. For calcium nitrate Ca(NO₃)₂: Atomic masses are H = 1 u, N = 14 u, O = 16 u, Ca = 40 u. Formula unit mass calculation: (40 × 1) + {(14 × 1) + (16 × 3)} × 2 = 40 + (14 + 48) × 2 = 40 + 62 × 2 = 40 + 124 = 164 u. Important note: Molecular mass applies to covalent materials (separate molecules). Formula unit mass applies to ionic materials (crystal lattice arrangement - no separate molecules). Nitric acid is covalent; calcium nitrate is ionic.
In simple words: Multiply each atom's mass by how many atoms are present, add them up. Use molecular mass for covalent compounds, formula unit mass for ionic compounds.

Exam Tip: Be careful with polyatomic ions - calculate the mass of the entire group first, then multiply by its subscript.

Question 10. 4.0 g of calcium carbonate reacts with 2.92 g of hydrochloric acid in a closed container to produce 1.76 g of CO₂, 0.72 g of water, and 4.44 g of calcium chloride. Verify the Law of Conservation of Mass.
Answer: Total mass of starting materials: Mass of CaCO₃ plus Mass of HCl equals 4.0 g plus 2.92 g equals 6.92 g. Total mass of products: Mass of CO₂ plus Mass of H₂O plus Mass of CaCl₂ equals 1.76 g plus 0.72 g plus 4.44 g equals 6.92 g. Since mass of starting materials (6.92 g) matches mass of products (6.92 g), the Law of Conservation of Mass is demonstrated and followed in this chemical process. This confirmation holds true since the process occurred inside a closed container - all gaseous products (CO₂) were maintained and weighed. If the arrangement were open, CO₂ would have left, making it look like the law was broken.
In simple words: Add reactants: 4.0 + 2.92 = 6.92 g. Add products: 1.76 + 0.72 + 4.44 = 6.92 g. They match, so the law is proven correct.

Exam Tip: Show your addition clearly on both sides of the equation and state that the law is verified when they are equal.

Long Answer Type Questions

Question 1. State the Law of Conservation of Mass. Describe Activity 9.2 to demonstrate it. Why is a closed system essential?
Answer: Law of Conservation of Mass: Matter cannot be generated or eliminated in a chemical process. Mass of starting materials equals mass of final products. Activity 9.2 works with vinegar and baking soda. In Set-up 1 (open arrangement), baking soda pours directly into vinegar held in an uncovered flask. CO₂ gas releases into the atmosphere. The final measured weight falls short of the starting weight - falsely appearing to defy the law. In Set-up 2 (sealed arrangement), a balloon filled with baking soda hangs from the flask's opening. As baking soda integrates with vinegar, CO₂ fills the balloon yet continues to exist within the arrangement. Final measured weight matches starting measured weight. A sealed arrangement proves essential since when any product (gas) releases outside, it gets unaccounted for in the final measurement - providing a misleading result that suggests weight was removed. The regulation stays universally accurate only when all starting points and results remain considered and weighed.
In simple words: Open system: gas escapes and is not weighed, so mass appears to decrease. Closed system: gas stays trapped, so total mass remains constant. The law always holds true in a closed system.

Exam Tip: Always emphasize the importance of a closed system - this is the key reason why mass appears to decrease in an open setup.

Question 2. State Dalton's Atomic Theory. How does it explain the Law of Conservation of Mass and the Law of Constant Proportions? What are its limitations?
Answer: Dalton's Atomic Theory (created in 1808) includes these core ideas: (i) All materials consist of minuscule atoms; (ii) atoms cannot be separated and cannot be generated or removed; (iii) atoms of a specific element match each other; (iv) atoms of distinct elements hold distinct masses; (v) atoms link in ratios of basic whole numbers to form compounds; (vi) a given compound maintains a fixed relative count and arrangement of atoms. The Law of Conservation of Mass is clarified since atoms only reorganize - no atom gets generated or removed, thus keeping complete mass steady. The Law of Constant Proportions is clarified since atoms forever link in the same fixed whole number proportions - producing a consistent mass proportion in any compound. Restrictions: Science later determined atoms are truly separable (include electrons, protons, neutrons). Atoms of an identical element can hold distinct masses (isotopes). These revelations post-dated Dalton, needing theory modifications.
In simple words: Atoms rearrange but do not change amount, so mass is conserved. Atoms always mix in the same way, so proportions stay constant. Later we found atoms can be divided.

Exam Tip: Know the full six postulates and be ready to explain both strengths (explains two laws) and weaknesses (atoms are divisible, isotopes exist).

Question 3. Explain the formation of a hydrogen chloride (HCl) molecule. Compare single and double covalent bonds with examples. How are covalent compounds named?
Answer: Hydrogen has atomic number 1 with one electron in K-shell needing one more for a consistent duplet. Chlorine carries atomic number 17 with setup 2, 8, 7 - with 7 valence electrons requiring one more for a consistent octet. Since both atoms require exactly one electron, each transfers one electron to the other. One shared electron pair creates the HCl molecule (single covalent bond). HCl functions as a covalent compound. Single bond: One shared electron pair. Examples: H₂ (H-H), Cl₂ (Cl-Cl), HCl (H-Cl). Double bond: Two shared electron pairs. Example: O₂ (O=O) - oxygen carries 6 valence electrons needing 2 additional, so two O atoms share 2 electrons each. Naming covalent compounds: Apply IUPAC prefix method - mono (1), di (2), tri (3), tetra (4), penta (5), hexa (6). Initial element maintains its name; subsequent element ends in -ide. Mono omitted for the initial element. Example: CO equals carbon monoxide; CO₂ equals carbon dioxide; PCl₃ equals phosphorus trichloride.
In simple words: Single bond - one pair shared. Double bond - two pairs shared. Name using prefixes and -ide ending for the second element.

Exam Tip: Be able to draw electron dot structures for both single and double covalent bonds, and practice naming various compounds using IUPAC rules.

Question 4. Explain the formation of an ionic bond using the example of NaCl. Describe the crystal structure of NaCl. Why does NaCl conduct electricity in solution but not in solid state?
Answer: Sodium (atomic number 11, setup 2, 8, 1) contains 1 valence electron. It readily transfers this electron to achieve the balanced neon setup 2, 8. Upon transfer, it converts to Na⁺ (11 protons, 10 electrons, charge +1). Chlorine (atomic number 17, setup 2, 8, 7) holds 7 valence electrons. It takes sodium's electron to achieve the balanced argon setup 2, 8, 8. Upon acceptance, it converts to Cl⁻ (17 protons, 18 electrons, charge -1). Na⁺ and Cl⁻ stick together by electrostatic pull - this is the ionic bond. They create NaCl and show no net charge (total charge = 0). Crystal structure: Ionic compounds do not create single molecules. In solid NaCl, every Na⁺ rests next to 6 Cl⁻ ions, and every Cl⁻ rests next to 6 Na⁺ ions, creating a regular 3-D crystal network pattern. Conductivity: In solid condition, all ions stay locked in position through intense electrostatic attractions - they cannot travel, so conductivity fails. When mixed in water, the crystal network breaks apart and Na⁺ and Cl⁻ ions become free to roam independently in the fluid. These roaming charged ions push electrical current, permitting conductivity.
In simple words: Solid NaCl has ions locked in place, so no conductivity. When dissolved, ions move freely in water, so it conducts electricity.

Exam Tip: Remember that ionic compounds must be in solution or molten state to conduct electricity because the ions must be able to move.

Question 5. Explain the criss-cross method for writing chemical formulae. Write the formulae for (a) magnesium hydroxide, (b) aluminium sulfate, (c) ferric chloride. Calculate the formula unit mass of magnesium hydroxide.
Answer: Criss-cross method for ionic compounds: First place cation symbol, followed by anion symbol. Write charge (numbers only) below each symbol. Exchange the charge digits as subscripts. Decrease subscripts by common factor when feasible. Apply brackets around polyatomic ions when subscript is more than 1. (a) Magnesium hydroxide Mg(OH)₂: Mg²⁺ and OH⁻ - criss-cross charges (2 and 1) yielding Mg(OH)₂. Brackets remain essential as subscript 2 affects the complete OH group. (b) Aluminium sulfate Al₂(SO₄)₃: Al³⁺ and SO₄²⁻ - criss-cross (3 and 2) yielding Al₂(SO₄)₃. No decrease feasible (2 and 3 show no common factor besides 1). (c) Ferric chloride FeCl₃: Fe³⁺ and Cl⁻ - criss-cross (3 and 1) yielding FeCl₃. Formula Unit Mass of Mg(OH)₂: Atomic masses: Mg = 24 u; O = 16 u; H = 1 u. Calculation: (24 × 1) + {(16 × 1) + (1 × 1)} × 2 = 24 + (17 × 2) = 24 + 34 = 58 u.
In simple words: Write ions, cross the charge numbers as subscripts, reduce by common factors, and use brackets for polyatomic ions with subscripts greater than 1.

Exam Tip: Practice the criss-cross method with various ions until it becomes automatic. Always check if subscripts can be reduced by a common factor.

Section-wise Notes and Summary

9.1 Law of Conservation of Mass

Demonstrated through Activities 9.1, 9.2 and 9.3:

Activity 9.1 (Physical change): Salt mass added to water mass produces salt solution mass. Mass stays unchanged through dissolution.
Activity 9.2 (Chemical change): Vinegar and baking soda react together.

• Experimental set-up 1: Gas leaves the system - mass appears to drop (measurement error)
• Experimental set-up 2: Balloon connected to flask captures gas - starting mass matches final mass

Activity 9.3 (Verification): Sodium sulfate plus Barium chloride produces Barium sulfate (white solid) plus Sodium chloride. Mass before equals mass after.
Law of Conservation of Mass: Matter cannot be made or removed in a chemical process. Complete mass of starting materials equals complete mass of final products.
Proposed by Antoine Lavoisier in 1789 - recognized as the founder of contemporary chemistry.
Key point: This law applies only when the arrangement is sealed (no gas leaves). In an open arrangement, apparent mass shifts happen as gas leaves - yet the complete mass including the missing gas continues to be stable.

9.2 Law of Constant Proportions

Suggested by Joseph Louis Proust - also referred to as the Law of Definite Proportions or Proust's Law. Statement: In any compound made from two or more elements, the components combine in a fixed mass proportion, independent of the source of the compound.
Example - Water: From a river, borehole, sea or any location - filtered water consistently includes hydrogen and oxygen in the mass ratio of 1:8. 9 g of water constantly releases 1 g of hydrogen and 8 g of oxygen when broken down.
Key distinction:

• Law of Constant Proportions applies to compounds (fixed mass ratio)
• It does NOT apply to mixtures (which can be combined in any ratio)

Both rules together establish the experimental base of Dalton's Atomic Theory.

9.3 Dalton's Atomic Theory

John Dalton (England, 1808) put forth his framework as a collection of concepts:

• Everything consists of minuscule bits called atoms, which engage in chemical processes.
• Atoms cannot be separated - they cannot be generated or eliminated in a chemical process.
• Atoms of a single element match in mass and chemical properties.
• Atoms of separate elements show distinct masses and chemical properties.
• Atoms join in the proportion of basic whole figures to create compounds.
• The fixed types and amounts of atoms are permanent in a given compound.

How Dalton's framework explains the two rules:

• Law of Conservation of Mass: Atoms reorganize only (never generated or eliminated) during a chemical process - overall mass stays stable.
• Law of Constant Proportions: Fixed kinds and numbers of atoms in a given compound - fixed proportion by mass.

Note: Later revelations (electrons, protons, neutrons) demonstrated atoms can be separated - Dalton's second concept needed updating. His framework provided the opening scientific explanation of how materials work in chemical processes.

9.4 How Atoms Combine?

A molecule is an electrically neutral unit made of multiple atoms that can remain by itself and exhibits all the traits of that material.
Note: Certain components like helium live only as atoms (not molecules) since their atoms hold already stable (complete valence shell).
Atoms link to stay stable by finishing an octet (or duplet for H and He). This occurs through two approaches:

1. Sharing of electrons - creates covalent bonds
2. Movement of electrons - creates ionic bonds

When atoms link, the total strength of the arrangement decreases - producing more stability. The attraction keeping atoms together is named a chemical bond.

9.4.1 Covalent Bond (Bonding by Sharing of Electrons)

A covalent bond happens when a pair of atoms distribute one or extra electron sets to establish balanced electron setups.
A. Molecules of Elements (identical element joining): Hydrogen molecule (H₂):

Covalent Bonding

A covalent bond forms when two atoms share one or more pairs of electrons. This shared pair of electrons is attracted to the nuclei of both atoms, holding them together.

Hydrogen molecule (H₂):

• H has 1 electron in K-shell; needs 1 more
• Two H atoms each share 1 electron → shared pair attracts both nuclei
• Single covalent bond (one shared pair) → represented as H—H

Chlorine molecule (Cl₂):

• Cl has 7 valence electrons; needs 1 more
• Two Cl atoms each share 1 electron → single bond → Cl—Cl

Oxygen molecule (O₂):

• O has 6 valence electrons; needs 2 more
• Two O atoms each share 2 electrons → double bond (two shared pairs) → O=O

Nitrogen molecule (N₂):

• N has 5 valence electrons; needs 3 more
• Two N atoms each share 3 electrons → triple bond → N≡N

Molecules of Compounds (different elements combining)

Hydrogen chloride (HCl):

• H needs 1 electron; Cl needs 1 electron
• Each shares 1 electron → single covalent bond → H—Cl

Water (H₂O):

• O needs 2 electrons; H needs 1 electron
• Two H atoms each share 1 electron with O → 2 single bonds → H—O—H
• Formula H₂O: 2 hydrogen atoms + 1 oxygen atom

Carbon dioxide (CO₂):

• C has 4 valence electrons; O has 6 valence electrons
• C forms double bond with each O → O=C=O

Naming Covalent Compounds (IUPAC Prefix System)

Rules:

• First element retains its regular name
• Second element ends in -ide
• If prefix ends in 'o' or 'a' and element starts with vowel, drop the last vowel (e.g., monoxide not monooxide)
• When hydrogen is the first element, no prefix is used for hydrogen (e.g., H₂S = hydrogen sulfide, not dihydrogen sulfide)

Examples:

• CO = carbon monoxide
• CO₂ = carbon dioxide
• CS₂ = carbon disulfide
• PCl₃ = phosphorus trichloride
• SF₆ = sulfur hexafluoride
• N₂O₄ = dinitrogen tetroxide
• N₂O₅ = dinitrogen pentoxide
• H₂O = water (common name)
• NH₃ = ammonia (common name)

Ionic Bond (Bonding by Electron Transfer)

An ionic bond forms when one atom transfers electrons to another, resulting in the formation of oppositely charged ions held together by electrostatic attraction.

Formation of sodium chloride (NaCl):

Sodium (Z=11, configuration 2, 8, 1):

• Has 1 valence electron → loses it → becomes Na⁺ (sodium cation)
• Na⁺ has 11 protons and 10 electrons → net charge = +1

Chlorine (Z=17, configuration 2, 8, 7):

• Has 7 valence electrons → gains 1 → becomes Cl⁻ (chloride anion)
• Cl⁻ has 17 protons and 18 electrons → net charge = -1

Na⁺ and Cl⁻ are held together by electrostatic force of attraction → ionic bond → NaCl

Key terms:

• Cation: Positive ion formed by loss of electrons (e.g., Na⁺, Ca²⁺, Al³⁺)
• Anion: Negative ion formed by gain of electrons (e.g., Cl⁻, O²⁻, S²⁻)
• Ions: Collective term for cations and anions
• Polyatomic ions: Ions formed by combination of two or more elements (e.g., SO₄²⁻, NO₃⁻, OH⁻, NH₄⁺)
• Ionic crystal structure: Ionic compounds form 3-D crystals. In NaCl, each Na⁺ is surrounded by 6 Cl⁻ ions, and each Cl⁻ is surrounded by 6 Na⁺ ions. This regular repeating pattern is called the crystal lattice.

Naming Ionic Compounds:

• Cation name is written first, anion name second
• Simple anions end in -ide (e.g., chloride, oxide, sulfide)
• Polyatomic anions do NOT end in -ide (e.g., sulfate, nitrate, carbonate)

Examples:

• NaCl = sodium chloride
• CaCO₃ = calcium carbonate
• Mg(OH)₂ = magnesium hydroxide

Writing Chemical Formulae

Covalent Compounds - Criss-Cross Method:

1. Write symbols of elements
2. Write their valencies below
3. Criss-cross the valencies as subscripts

Examples:

• HCl: H(1) and Cl(1) → criss-cross → HCl (subscript 1 is not written)
• H₂S: H(1) and S(2) → criss-cross → H₂S
• CCl₄: C(4) and Cl(1) → criss-cross → CCl₄

Ionic Compounds - Criss-Cross Method:

1. Write cation symbol first, then anion
2. Write charges (numbers only) below
3. Criss-cross the charge numbers as subscripts
4. Reduce subscripts by common factor if needed
5. Use brackets for polyatomic ions when subscript > 1

Examples:

• CaCl₂: Ca²⁺ and Cl⁻ → criss-cross → CaCl₂
• Al₂O₃: Al³⁺ and O²⁻ → criss-cross → Al₂O₃
• MgO: Mg²⁺ and O²⁻ → criss-cross gives Mg₂O₂ → reduce to MgO
• CaCO₃: Ca²⁺ and CO₃²⁻ → same valency → reduce to CaCO₃
• Mg(OH)₂: Mg²⁺ and OH⁻ → criss-cross → Mg(OH)₂ (brackets needed)
• Al(OH)₃: Al³⁺ and OH⁻ → criss-cross → Al(OH)₃ (NOT AlOH₃)
• Al₂(SO₄)₃: Al³⁺ and SO₄²⁻ → criss-cross → Al₂(SO₄)₃

Note: Charges on ions are NOT indicated in the formula of the compound.

Properties of Ionic and Covalent Compounds

Key explanation: Ionic compounds conduct electricity only when dissolved in water or in molten state because ions become free to move. In solid state, ions are held rigidly in crystal lattice and cannot move. Sugar is a covalent compound that dissolves in water but does NOT conduct electricity because it does not produce ions in solution.

Molecular Mass of Covalent Compounds

Molecular mass represents the sum of atomic masses of all atoms present in one molecule.

Formula: Molecular mass = (atomic mass of each element × number of atoms of that element), sum for all elements

Examples:

• Water (H₂O): (1×2) + (16×1) = 18 u
• Carbon dioxide (CO₂): (12×1) + (16×2) = 44 u
• Methane (CH₄): (12×1) + (1×4) = 16 u
• Nitric acid (HNO₃): (1×1) + (14×1) + (16×3) = 63 u

Note: Ionic compounds do NOT have molecular mass because they form 3-D crystal structures, not discrete molecules.

Formula Unit Mass of Ionic Compounds

A formula unit represents the collection of the simplest whole number ratio of ions in an ionic compound. Formula unit mass equals the sum of atomic masses of all atoms in a formula unit.

Examples:

• Na₂O: (23×2) + (16×1) = 62 u
• Ca(NO₃)₂: (40×1) + {(14×1) + (16×3)} × 2 = 40 + 62×2 = 164 u
• KCl: (39×1) + (35.5×1) = 74.5 u
• Mg(OH)₂: (24×1) + {(16×1) + (1×1)} × 2 = 24 + 34 = 58 u

Important Formulae and Key Notes

Common Atomic Masses for Calculations: H = 1 u, C = 12 u, N = 14 u, O = 16 u, Na = 23 u, Mg = 24 u, Al = 27 u, S = 32 u, Cl = 35.5 u, K = 39 u, Ca = 40 u, Fe = 56 u, Cu = 64 u, Zn = 65 u