An Insight Into Objective Chemistry-11, Ch-01

Here are complete topic-wise short questions and MCQs of our famous publication “An Insight Into Objective Chemistry-11”. Additional details and diagrams have been in this digital product, so you may call it plus or updated version of our hard form.

1.1: ATOM

1.1.1: Evidence of Atom

1.1.2: Molecule

1.1.3: Ion

1.1.4: Molecular Ion



1.3.1: Relative Abundance of Isotopes

1.3.2: Determination of Relative Atomic Masses of Isotopes by Mass Spectrometry

1.3.3: Average Atomic Masses


1.4.1: Empirical Formula

1.1.1: Empirical Formula From Combustion Analysis

1.1.1: Molecular Formula


1.1.1: Avogadro’s Number

1.1.1: Molar Volume



1.1.1: YIELD

Q.01: What is the contribution of Greeks to the study of atom?

Q.02: What is the significance of John Dalton’s work about atom?

Q.03: What is the contribution of J. Berzelius in chemistry?

Q.04: Why can atom not be seen with optical microscope?

Q.05: How can atom be seen with electron microscope?

Q.06: Differentiate between an ‘atom’ and a ‘molecule’.

Q.07: Define the term atomicity. Give examples.

Q.08: What are macromolecules? Give an example.

Q.09: What are ions? Under what conditions are they produced?

Q.10: The formation of a positive ion is an endothermic reaction. Explain.

Q.11: The formation of a uninegative ion is an exothermic reaction. Explain.

Q.12: Define molecular ions with examples.

Q.13: What is molecular ion? Give an example.

Q.14: What are molecular ions? How are they generated?

Q.15: Differentiate between ion and molecular ion.

Q.16: What is relative atomic mass?

Q.17: Why we use the term relative atomic mass?

Q.18: What is amu? 

Q.19: Define isotopes and isotopy.

Q.20: What are isotopes? Give an example.

Q.21: Why do isotopes have same chemical but different physical properties?

Q.22: What are mono-isotopic elements?

Q.23: Define mass spectrometer and mass spectrometry.

Q.24: What is the principle of mass spectrometry?

Q.25: Define mass spectrum.

Q.26: Draw a labelled diagram of Dempster’s mass spectrometer.

Q.27: What is the function of ionization chamber in mass spectrometer?

Q.28: What is the function of magnetic field in mass spectrometer?

Q.29: What is electrometer? Give its function in mass spectrometer.

Q.30: Write down any four methods used for the separation of isotopes.

Q.31: Why do most of the elements have fractional atomic masses?

Q.32: No individual Neon atom in the sample of the element has a mass of 20.18 amu. Explain.

Q.33: What is the justification of two equally strong peaks in the mass spectrum for bromine, while for iodine only one peak at 127amu is indicated?

Q.34: Differentiate between qualitative and quantitative analyses.

Q.35: Write down the formulas to determine the percentage of carbon and hydrogen in a compound.

Q.36: Calculate the percentage of nitrogen in NH2CONH2.

Q.37: What are the steps to determine the empirical formula? 

Q.38: Why can the percentage of oxygen not be calculated directly in combustion analysis?             

Q.39: Differentiate between empirical formula and molecular formula.

Q.40: Define Empirical Formula mass.

Q.41: Define molecular formula. How is it related to empirical formula?

Q.42: How can you justify with example that molecular formula = n ´ empirical formula?

Q.43: A compound may have same empirical and molecular formulas. Justify giving examples.

Q.44: Define mole with two examples.

Q.45: Define gram atom, giving examples.

Q.46: Define gram molecule.

Q.47: Define gram atom and gram molecule.

Q.48: Define gram formula. Give two examples.

Q.49: Define gram ion.                                                                                                    

Q.50: Differentiate between gram atom and gram ion.

Q.51: Define molar mass.

Q.52: Mg atom is twice heavier as an atom of carbon. Explain.

Q.53: Justify that N2 and CO have same number of electrons, protons and neutrons.

Q.54: One mg of K2CrO4 has thrice the number of ions than the number of formula units when ionized in water. Justify.

Q.55: 23 g of Na and 238 g of Uranium have equal number of atoms in them. Justify it.

Q.56: 180 g of glucose and 342 g of sucrose have the same number of molecules but different number of atoms present in them. Give reason.

Q.57: 58.5 amu is the formula mass of NaCl, but not the molecular mass. Why?

Q.58: Calculate the number of gram atoms (moles) in 0.1 g of sodium.

Q.59: Calculate the number of gram atoms in 0.1 kg of silicon.

Q.60: How many molecules of water are there in 10 g of ice?

Q.61: How many molecules of water are present in 3.6 grams of water?

Q:62: How many moles are present in 18 g of H2O?

Q.63: Calculate the mass in grams of 10–3 moles of water.

Q.64: Calculate the mass in grams of 2.74 moles of KMnO4.

Q.65: Calculate the mass in kilograms of 2.6 × 1020 molecules of SO2.

Q.66: Calculate the moles of chlorine atoms in 0.822 g of C2H4Cl2.

Q.67: Calculate mass in grams of 5.136 moles of Ag2CO3.

Q.68: Calculate the mass in grams of 2.78 ´ 1021 molecules of CrO2Cl2.

Q.69: Define mole and Avogadro’s number.

Q.70: Define Avogadro’s number. Give an equation to relate the Avogadro’s number and the mass of an element.

Q.71: When H2SO4 is dissolved in water, the number of ions is unequal, but the number of charges is equal. Explain.

Q.72: Define molar volume.

Q.73: One dm3 of different gases H2 and O2 have different masses but occupy same volume. Explain.

Q.74: Define Stoichiometry. Give two assumptions for stoichiometric calculations. OR Define stoichiometry. Write down its two laws.

Q.75: What are the limitations of a chemical equation?

Q.76: Law of conservation of mass has to be obeyed during stoichiometric calculations. Explain.

Q.77: One mole H2SO4 requires 2 moles of NaOH to neutralize it. Explain.

Q.78: How many moles of CO2 can be produced from burning one mole of octane? Molar mass of octane is 114.

Q.79: Define limiting reactant. Give an example.

Q.80: Limiting reactant controls the amount of the product. Explain.

Q.81: Why in some reactions, one of the reactants is used deliberately in excess quantity?

Q.82: What are the steps to identify a limiting reactant?

Q.83: Many chemical reactions taking place in our surroundings involve limiting reactants. Explain.

Q.84: Differentiate between actual yield and theoretical yield.

Q.85: Why in most reactions, actual yield remains less than the theoretical yield?

Q.86: Define percentage yield. Give an example.

Q.87: What is the percentage yield?

Q.88: How can the efficiency of chemical reaction be calculated? Or What is the percentage yield?

01: Explain isotopes with their relative abundance. (BWP-22)

02: Explain the construction and working of mass spectrometer. (RWP-22)

03: What is empirical formula? Discuss steps to calculate empirical formula. (BWP-23)

04: Write down various steps to calculate the empirical formula of a compound. (FSL-22)

05: Describe combustion analysis to determine the percentage of C, H and O in the organic compound. (LHR-22)(FSL-22)(DGK-22)(SRG-23)

06: Explain combustion analysis with diagram and write formulas for percentage of carbon, hydrogen and oxygen. (GJR-23)

07: A sample of liquid consisting of carbon, hydrogen and oxygen was subjected to combustion analysis. 0.5439 g of the compound gave 1.039 g of CO2, 0.6369 g of H2O. Determine the empirical formula of the compound. (MLT-17)(DGK-16,16,18)

08: The combustion analysis of an organic compound shows it to contain 65.44% carbon, 5.50% hydrogen and 29.06% or oxygen. What is the empirical formula of the compound? If the molecular mass of this compound is 110.15 g mol-1. Calculate the molecular formula of the compound. (RWP-16)(LHR-16)                             

09: Ethylene glycol is used as automobile antifreeze. It has 38.7% carbon, 9.7% hydrogen and 51.6% oxygen. Its molar mass is 62.1 g mole-1. Determine its empirical formula. (LHR-17) 

10: A well-known ideal gas is enclosed in a container having volume 500 cm3 at S.T.P. Its mass comes out to be 0.72 g. What is the molar mass of this gas? (AJK-16,17)(MLT-16,16,17)(DGK-17)(BWP-18)(SWL-18)

11: Define the following terms with two, two examples of each: 1) Gram Formula 2) Gram Ion 3) Gram Atom 4) Percentage Yield (DGK-22)

12: Define the following terms with examples: (i) Relative atomic mass (ii) Molecular ion (iii) Isotope (iv) Molar volume (LHR-23)

13: Calculate the gram atoms (moles) in: (i) 0.1 g of sodium (ii) 0.1 kg of silicon. (DGK-21)

14: 10 grams of H3PO4 have been dissolved in excess of water to dissolve completely into ions. Calculate: (i) Masses of individual ions. (ii) Number of positive and negative charges dispersed in solution. (MLT-21)            

15: Define stoichiometry. Write down its assumptions. (SWL-23)

16: Define stoichiometry. Give its assumptions. Mention two laws which help to perform the stoichiometric calculations. (BWP-23)(MLT-23)(RWP-23)

17: Mg metal reacts with HCl to give hydrogen gas. What is the minimum volume of HCl solution (27% by weight) required to produce 12.1 g of H2? The density of HCl solution is 1.14 g/cm3. (DGK-17,21)(MLT-21)

18: Calculate the number of grams of K2SO4 and water produced when 14 g of KOH are reacted with excess of H2SO4. Also calculate the number of molecules of water produced? (LHR-17,21)(DGK-21)(FSL-21)(FSL-21)

19: Serotonin (Molar mass = 176 g mol–1) is a compound that conducts nerve impulses in brain and muscles. It contains 68.2% C, 6.86% H, 15.09% N, and 9.08% O. What is its molecular formula? (BWP-17)                                                                 

20: Write a note on limiting reactant and explain it with help of two examples. (SWL-22)(RWP-22)(FSL-23)

21: What is limiting reactant? How does it control the quantity of product in a chemical reaction? Give two examples. (BWP-22)

22: What is limiting reactant? Give two examples. How limiting reactant can be identified? (GJR-23)

23: Define limiting reactant. And give the suitable examples, and also write down the steps to find limiting reactant. (GJR-22)(SWL-23)

24: NH3 gas can be prepared by heating together two solids NH4Cl and Ca(OH)2. If a mixture containing 100 grams of each solid is heated, then how many grams of NH3 is produced? 2NH4Cl(s) + Ca(OH)2 (s) ⟶ CaCl2(s) +2NH3(g) + 2H2O(ℓ) (LHR-16,21)(MLT-16)(BWP-17)(RWP-21)

25: Define yield. Differentiate between actual and theoretical yield. How percentage yield can be calculated? (DGK-23)

26: What is the difference between actual yield and theoretical yield? Why actual yield is less than theoretical yield? (RWP-22)(GJR-22)

27: Define types of yield. How do we calculate the percentage yield of a reaction? (LHR-23)

28: When lime stone (CaCO3) is roasted, quick lime (CaO) is produced according to the following equation. The actual yield of CaO is 2.5 kg, when 4.5 kg of lime stone is roasted. What is the percentage yield of this reaction? CaCO3 (s) ⟶ CaO(s) +CO2(g) (FSL-16,17,21)(SRG-17)(RWP-17,17)(GJR-21)

29: Silicon carbide (SiC) is an important ceramic material. It is produced by allowing sand (SiO2) to react with carbon at high temperature.  SiO2 + 3C ⟶ SiC + 2CO When 100 kg sand is reacted with excess of carbon, 51.4 kg of SiC is produced. What is the percentage yield of SiC?(DGK-18)

Textbook Exercise

Previous Boards Essentials

Textbook Conceptuals

1.1: Atom

1.2: Relative Atomic Mass

1.3: Isotopes

1.4: Analysis of a Compound; Empirical and Molecular Formulas

1.5: Concept of Mole

1.6: Stoichiometry

1.7: Limiting Reactant

1.8: Yield

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