11 chemistry assignment answer key 2021

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Tamilnadu 11th Time Table | Plus One Exam Time Table

Our Tamilnadu School Education Department conducts the public exam for 11th standard every year. Due to corona the previous public exams were cancelled by Tamilnadu Government. Every year the public exam will be conducted in March to April, But this year it is planned to conduct in the month of may 2022.

In our website, we published so many study materials, model question papers, with proper answer keys. We published lot of materials on the basis of the reduced sylabus. If you have any problem in downloading any materials, you can comment below and if you want to upload any materials from your side in order to help the students, you can mail your materials to [email protected]

Bookmark this page using the keyboard shortcut Ctrl + D to refer to this page and check your results and other exam-related updates. All the Best!

11th Chemistry - Public Exam 2021-2022 - Official Answer Keys | DGE - (English Medium) - CLICK TO DOWNLOAD

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• NCERT Solutions
• NCERT Class 11
• NCERT 11 Chemistry
• Chapter 1: Some Basic Concepts Chemistry

NCERT Solutions for Class 11 Chemistry Chapter 1: Some Basic Concepts of Chemistry

Ncert solutions class 11 chemistry chapter 1 – free pdf download.

NCERT Solutions for Class 11 Chemistry Chapter 1 Some Basic Concepts of Chemistry are provided on this page for Class 11 Chemistry students. These solutions can also be downloaded in a PDF format for free by clicking the download button provided below. The free Class 11 Chemistry NCERT Solutions provided by BYJU’S have been prepared by seasoned academics and chemistry experts, keeping the needs of Class 11 students in mind. The NCERT Solutions provided on this page contain step-by-step explanations in order to help students answer similar questions that might be asked in their board examinations.

The expert tutors have created NCERT Solutions for Class 11 in simple and understandable language in accordance with the latest CBSE Syllabus 2023-24 and its guidelines. The basic concepts are covered in the NCERT Class 11 Solutions with the aim of providing a quality learning experience for Class 11 students. You can download the Class 11 Chemistry NCERT Solutions Chapter 1 PDF from BYJU’S to complete a part of the CBSE syllabus before the board exam.

Download NCERT Solutions Class 11 Chemistry Chapter 1 PDF

“Some Basic Concepts of Chemistry” is the first chapter in the Class 11 Chemistry syllabus as prescribed by NCERT . The chapter touches upon topics such as the importance of Chemistry, atomic mass, and molecular mass. Some basic laws and theories in Chemistry, such as Dalton’s atomic theory, Avogadro’s law and the law of conservation of mass, are also discussed in this chapter.

The types of questions provided in the NCERT Class 11 Chemistry textbook for Chapter 1 include:

• Numerical problems in calculating the molecular weight of compounds.
• Numerical problems in calculating mass percent and concentration.
• Problems on empirical and molecular formulae.
• Problems on molarity and molality.
• Other problems related to the mole concept (such as percentage composition and expressing concentration in parts per million).

The NCERT Solutions of Chemistry provided on this page for Class 11 Chapter 1 contains detailed explanations of the steps to be followed while solving the numerical value questions that are frequently asked in the examinations. The subtopics covered in the chapter are listed below.

Subtopics of NCERT Solutions for Class 11 Chemistry Chapter 1 Some Basic Concepts of Chemistry

• Importance of Chemistry
• Nature of Matter
• Properties of Matter and Their Measurement  
• The International System of Units (SI)
• Mass and Weight
• Uncertainty in Measurement
• Scientific Notation
• Significant Figures
• Dimensional Analysis
• Laws of Chemical Combinations    
• Law of Conservation Of Mass    
• Law of Definite Proportions
• Law of Multiple Proportions  
• Gay Lussac’s Law of Gaseous Volumes  
• Avogadro’s Law
• Dalton’s Atomic Theory
• Atomic and Molecular Masses   
• Atomic Mass
• Average Atomic Mass
• Molecular Mass
• Formula Mass
• Mole Concept and Molar Masses
• Percentage Composition    
• Empirical Formula for Molecular Formula
• Stoichiometry and Stoichiometric Calculations                
• Limiting Reagent
• Reactions in Solutions

Access the answers of NCERT Solutions for Class 11 Chemistry Chapter 1

Q1. Calculate the molar mass of the following:

(i) \(\begin{array}{l}CH_{4}\end{array} \)       (ii) \(\begin{array}{l}H_{2}O\end{array} \)       (iii) \(\begin{array}{l}CO_{2}\end{array} \)

(i)  \(\begin{array}{l}CH_{4}\end{array} \) :

Molecular mass of \(\begin{array}{l}CH_{4}\end{array} \) = Atomic mass of C + 4 x Atomic mass of H

= 12 + 4 x 1

(ii) \(\begin{array}{l}H_{2}O\end{array} \) :

Molar mass of water \(\begin{array}{l}H_{2}O\end{array} \)

Atomic mass of H = 1

Atomic mass of O = 16

H 2 O = 2×H+1×O

Molar mass of water = 2×1+16 = 18g/mol

(iii) \(\begin{array}{l}CO_{2}\end{array} \) :

Molecular mass of \(\begin{array}{l}CO_{2}\end{array} \) = Atomic mass of C + 2 x Atomic mass of O

= 12 + 2 × 16

Q2. Calculate the mass per cent of different elements present in sodium sulphate ( \(\begin{array}{l}Na_{2}SO_{4}\end{array} \) ) .

Now for \(\begin{array}{l}Na_{2}SO_{4}\end{array} \) .

Molar mass of \(\begin{array}{l}Na_{2}SO_{4}\end{array} \)

= [(2 x 23.0) + (32.066) + 4(16.00)]

Formula to calculate mass percent of an element = \(\begin{array}{l}\frac{Mass\;of\;that\;element\;in\;the\;compound}{Molar\;mass\;of\;the\;compound}\times 100\end{array} \)

Therefore, mass percent of the sodium element:

= \(\begin{array}{l}\frac{46.0g}{142.066g}\times 100\end{array} \)

Mass percent of the sulphur element:

= \(\begin{array}{l}\frac{32.066g}{142.066g}\times 100\end{array} \)

Mass percent of the oxygen element:

= \(\begin{array}{l}\frac{64.0g}{142.066g}\times 100\end{array} \)

Q3. Determine the empirical formula of an oxide of iron, which has 69.9% iron and 30.1% dioxygen by mass.

Given there is an oxide of iron which has 69.9% iron and 30.1% dioxygen by mass:

Relative moles of iron in iron oxide:

= \(\begin{array}{l}\frac{percent\;of\;iron\;by\;mass}{Atomic\;mass\;of\;iron}\end{array} \)

= \(\begin{array}{l}\frac{69.9}{55.85}\end{array} \)

Relative moles of oxygen in iron oxide:

= \(\begin{array}{l}\frac{percent\;of\;oxygen\;by\;mass}{Atomic\;mass\;of\;oxygen}\end{array} \)

= \(\begin{array}{l}\frac{30.1}{16.00}\end{array} \)

The simplest molar ratio of iron to oxygen:

⇒  1.25: 1.88 ⇒  1: 1.5 ⇒  2: 3

Therefore, the empirical formula of the iron oxide is \(\begin{array}{l}Fe_{2}O_{3}\end{array} \) .

Q4. Calculate the amount of carbon dioxide that could be produced when (i) 1 mole of carbon is burnt in air. (ii) 1 mole of carbon is burnt in 16 g of dioxygen. (iii) 2 moles of carbon are burnt in 16 g of dioxygen.

(i) 1 mole of carbon is burnt in air.

1 mole of carbon reacts with 1 mole of O 2 to form one mole of CO 2 .

Amount of \(\begin{array}{l}CO_{2}\end{array} \) produced = 44 g

(ii) 1 mole of carbon is burnt in 16 g of O 2 .

1 mole of carbon burnt in 32 grams of O 2 it forms 44 grams of \(\begin{array}{l}CO_{2}\end{array} \) .

Therefore, 16 grams of O 2 will form \(\begin{array}{l}\frac{44\times 16}{32}\end{array} \)

= 22 grams of \(\begin{array}{l}CO_{2}\end{array} \)

(iii) 2 moles of carbon are burnt in 16 g of O 2 .

Here again, dioxygen is the limiting reactant. 16g of dioxygen can combine only with 0.5mol of carbon. CO 2 produced again is equal to 22g.

Q5. Calculate the mass of sodium acetate \(\begin{array}{l}(CH_{3}COONa)\end{array} \) required to make 500 mL of 0.375 molar aqueous solution. Molar mass of sodium acetate is 82.0245 g mol –1 .

0.375 M aqueous solution of \(\begin{array}{l}CH_{3}COONa\end{array} \)

= 1000 mL of solution containing 0.375 moles of \(\begin{array}{l}CH_{3}COONa\end{array} \)

Therefore, no. of moles of \(\begin{array}{l}CH_{3}COONa\end{array} \) in 500 mL

= \(\begin{array}{l}\frac{0.375}{1000}\times 500\end{array} \)

= 0.1875 mole

Molar mass of sodium acetate = \(\begin{array}{l}82.0245\;g\;mol^{-1}\end{array} \)

Therefore, the mass of \(\begin{array}{l}CH_{3}COONa\end{array} \)

= \(\begin{array}{l}(82.0245\;g\;mol^{-1})(0.1875\;mole)\end{array} \)

= 15.38 grams

Q6. Calculate the concentration of nitric acid in moles per litre in a sample which has a density, 1.41 g mL –1 and the mass per cent of nitric acid in it being 69%.

Mass percent of HNO 3 in sample is 69 %

Thus, 100 g of HNO 3 contains 69 g of HNO 3 by mass.

Molar mass of HNO 3

= { 1 + 14 + 3(16)} \(\begin{array}{l}g.mol^{-1}\end{array} \)

= 1 + 14 + 48

Now, no. of moles in 69 g of \(\begin{array}{l}HNO_{3}\end{array} \) :

= \(\begin{array}{l}\frac{69\:g}{63\:g\:mol^{-1}}\end{array} \)

= 1.095 mol

Volume of 100g HNO 3 solution

= \(\begin{array}{l}\frac{Mass\;of\;solution}{density\;of\;solution}\end{array} \)

= \(\begin{array}{l}\frac{100g}{1.41g\;mL^{-1}}\end{array} \)

= \(\begin{array}{l}70.92\times 10^{-3}\;L\end{array} \)

Concentration of HNO 3

= \(\begin{array}{l}\frac{1.095\:mole}{70.92\times 10^{-3}L}\end{array} \)

= 15.44mol/L

Concentration of HNO 3 = 15.44 mol/L

Q7. How much copper can be obtained from 100 g of copper sulphate (CuSO 4 )?

1 mole of \(\begin{array}{l}CuSO_{4}\end{array} \) contains 1 mole of Cu.

Molar mass of \(\begin{array}{l}CuSO_{4}\end{array} \)

= (63.5) + (32.00) + 4(16.00)

= 63.5 + 32.00 + 64.00

= 159.5 grams

159.5 grams of \(\begin{array}{l}CuSO_{4}\end{array} \) contains 63.5 grams of Cu.

Therefore, 100 grams of \(\begin{array}{l}CuSO_{4}\end{array} \) will contain \(\begin{array}{l}\frac{63.5\times 100g}{159.5}\end{array} \) of Cu.

= \(\begin{array}{l}\frac{63.5\times 100}{159.5}\end{array} \)

=39.81 grams

Q8. Determine the molecular formula of an oxide of iron, in which the mass percent of iron and oxygen are 69.9 and 30.1, respectively. 

Mass percent of Fe = 69.9%

Mass percent of O = 30.1%

No. of moles of Fe present in oxide

= \(\begin{array}{l}\frac{69.90}{55.85}\end{array} \)

No. of moles of O present in oxide

= \(\begin{array}{l}\frac{30.1}{16.0}\end{array} \)

Ratio of Fe to O in oxide,

= 1.25: 1.88

= \(\begin{array}{l}\frac{1.25}{1.25}:\frac{1.88}{1.25}\end{array} \)

= \(\begin{array}{l}1:1.5\end{array} \)

= \(\begin{array}{l}2:3\end{array} \)

Therefore, the empirical formula of oxide is \(\begin{array}{l}Fe_{2}O_{3}\end{array} \)

Empirical formula mass of \(\begin{array}{l}Fe_{2}O_{3}\end{array} \)

= [2(55.85) + 3(16.00)] g

The molar mass of  \(\begin{array}{l}Fe_{2}O_{3}\end{array} \) = 159.69g

Therefore n = \(\begin{array}{l}\frac{Molar\;mass}{Empirical\;formula\;mass}=\frac{159.69\;g}{159.7\;g}\end{array} \)

= 1(approx)

The molecular formula of a compound can be obtained by multiplying n with the empirical formula.

Thus, the empirical of the given oxide is \(\begin{array}{l}Fe_{2}O_{3}\end{array} \) and n is 1.

Therefore, the molecular formula of the oxide is  \(\begin{array}{l}Fe_{2}O_{3}\end{array} \)

Q9 . Calculate the atomic mass (average) of chlorine using the following data:

Fractional Abundance of 35 Cl = 0.7577 and Molar mass = 34.9689  

Fractional Abundance of 37 Cl = 0.2423 and Molar mass = 36.9659 

Average Atomic mass = (0.7577 x 34.9689)amu + (0.2423 x 36.9659) 

= 26.4959 + 8.9568 = 35.4527

Q10. In three moles of ethane (C 2 H 6 ), calculate the following: (i) Number of moles of carbon atoms. (ii) Number of moles of hydrogen atom

(iii) Number of molecules of ethane

(i) 1 mole of \(\begin{array}{l}C_{2}H_{6}\end{array} \) contains two moles of C- atoms.

(ii) 1 mole of \(\begin{array}{l}C_{2}H_{6}\end{array} \) contains six moles of H- atoms.

(iii) 1 mole of \(\begin{array}{l}C_{2}H_{6}\end{array} \) contains 1 mole of ethane- atoms.

= 3 x 6.023 x \(\begin{array}{l}10^{23}\end{array} \)

= 18.069 x \(\begin{array}{l}10^{23}\end{array} \)

Q11. What is the concentration of sugar (C 12 H 22 O 11 ) in mol L –1 if its 20 g are dissolved in enough water to make a final volume up to 2L?

Molarity (M) is as given by,

= \(\begin{array}{l}\frac{Number\;of\;moles\;of\;solute}{Volume\;of\;solution\;in\;Litres}\end{array} \)

= \(\begin{array}{l}\frac{\frac{Mass\;of\;sugar}{Molar\;mass\;of\;sugar}}{2\;L}\end{array} \)

= \(\begin{array}{l}\frac{\frac{20\;g}{[(12\;\times \;12)\;+\;(1\;\times \;22)\;+\;(11\;\times \;16)]g]}}{2\;L}\end{array} \)

= \(\begin{array}{l}\frac{\frac{20\;g}{342\;g}}{2\;L}\end{array} \)

= \(\begin{array}{l}\frac{0.0585\;mol}{2\;L}\end{array} \)

= 0.02925 mol \(\begin{array}{l}L^{-1}\end{array} \)

Therefore, Molar concentration = 0.02925 mol \(\begin{array}{l}L^{-1}\end{array} \)

Q12. If the density of methanol is 0.793 kg L –1 , what is its volume needed for making 2.5 L of its 0.25 M solution?

Molar mass of methanol (CH 3 OH)

= 32 gmol -1 = 0.032 kgmol -1

molarity of the given solution

24.78 x V 1 = 0.25 x 2.5 L

or V 1 = 0.02522L = 25.22mL

Q13. Pressure is determined as force per unit area of the surface. The SI unit of pressure, pascal is as shown below: 1Pa = 1N m –2 If mass of air at sea level is 1034 g cm –2 , calculate the pressure in pascal

Pressure is the force (i.e., weight) acting  per unit area

But weight = mg

∴ Pressure = Weight per unit area

Q14. What is the SI unit of mass? How is it defined?

The SI unit of mass is kilogram (kg). A kilogram is equal to the mass of a platinum-iridium cylinder kept at the International Bureau of Weights and Measures at Sèvres, France.

Q15. Match the following prefixes with their multiples:

Q16. What do you mean by significant figures?

Significant figures are the meaningful digits which are known with certainty. Significant figures indicate uncertainty in experimented value.

e.g.: The result of the experiment is 15.6 mL in that case 15 is certain and 6 is uncertain. The total significant figures are 3.

Therefore, “the total number of digits in a number with the last digit that shows the uncertainty of the result is known as significant figures.”

Q17. A sample of drinking water was found to be severely contaminated with chloroform, CHCl 3 , supposed to be carcinogenic in nature. The level of contamination was 15 ppm (by mass). (i) Express this in per cent by mass. (ii) Determine the molality of chloroform in the water sample.

(i) 1 ppm = 1 part out of 1 million parts.

Mass percent of 15 ppm chloroform in H 2 O

= \(\begin{array}{l}\frac{15}{10^{6}} \times 100\end{array} \)

= \(\begin{array}{l}\approx\end{array} \) 1.5 × \(\begin{array}{l}10^{-3}\end{array} \) %

(ii)  \(\begin{array}{l}Molarity = \frac{15/119.5}{10^{6}\times 10^{-3}}= 1.25 \times 10^{-4}\end{array} \)

Q18. Express the following in the scientific notation: (i) 0.0048 (ii) 234,000 (iii) 8008 (iv) 500.0 (v) 6.0012

(i) 0.0048= 4.8 × \(\begin{array}{l}10^{-3}\end{array} \)

(ii) 234,000 = 2.34 × \(\begin{array}{l}10^{5}\end{array} \)

(iii) 8008= 8.008 × \(\begin{array}{l}10^{3}\end{array} \)

(iv) 500.0 = 5.000 × \(\begin{array}{l}10^{2}\end{array} \)

(v) 6.0012 = 6.0012 × \(\begin{array}{l}10^{0}\end{array} \)

Q19. How many significant figures are present in the following?

(d) 126,000

(a) 0.0025: 2 significant numbers.

(b) 208: 3 significant numbers.

(c) 5005: 4 significant numbers.

(d) 126,000:3 significant numbers.

(e) 500.0: 4 significant numbers.

(f) 2.0034: 5 significant numbers.

Q20. Round up the following upto three significant figures:

(b) 10.4107

(a) The number after round up is: 34.2

(b) The number after round up is: 10.4

(c)The number after round up is: 0.0460

(d)The number after round up is: 2810

Q21. The following data are obtained when dinitrogen and dioxygen react together to form different compounds:

(a) Which law of chemical combination is obeyed by the above experimental data? Give its statement. (b) Fill in the blanks in the following conversions: (i) 1 km = …………………. mm = …………………. pm (ii) 1 mg = …………………. kg = …………………. ng (iii) 1 mL = …………………. L = …………………. dm 3

Here if we fix the mass of dinitrogen at 14g, then the masses of dioxygen that will combine with the fixed mass of dinitrogen are 16g, 32g, 32g, and 80g.

The masses of dioxygen bear a whole number ratio of 1:2:2:5.

Hence, the given experimental data obeys the Law of Multiple Proportions.

i.  \(\begin{array}{l}1 km = 1 km \times \frac{1000m}{1km}\times \frac{100cm}{1m}\times \frac{10mm}{1cm}= 10^{6}mm\\1km = 1 km \times \frac{1000m}{1km}\times \frac{1pm}{10^{-12}m}= 10^{^{15}}pm\end{array} \)

ii.  \(\begin{array}{l}1 mg = 1 mg\times \frac{1g}{1000mg}\times \frac{1kg}{1000g}= 10^{-6}kg\\1mg = 1mg\times \frac{1g}{1000mg}\times \frac{1ng}{10^{-9}g}=10^{6}ng\end{array} \)

iii.  \(\begin{array}{l}1mL = 1mL\times \frac{1L}{1000mL}=10^{-3}L\\1mL = 1cm^{3} \\ =1cm^{3}\times \frac{1dm\times 1dm\times 1dm}{10cm\times 10cm\times 10cm}= 10^{-3}dm^{3}\end{array} \)

Q22. If the speed of light is 3.0 × 10 8 m s –1 , calculate the distance covered by light in 2.00 ns.

Time taken = 2 ns

= 2 × \(\begin{array}{l}10^{ -9 }\end{array} \) s

Speed of light = 3 × \(\begin{array}{l}10^{ 8 }\end{array} \) \(\begin{array}{l}ms^{ -1 }\end{array} \)

We know that,

Distance = Speed x Time

Distance travelled in 2 ns = speed of light x time taken

= (3 × \(\begin{array}{l}10^{ 8 }\end{array} \) )(2 × \(\begin{array}{l}10^{ -9 }\end{array} \) )

= 6 × \(\begin{array}{l}10^{ -1 }\end{array} \) m

Q23. In a reaction A + B 2 →  AB 2 Identify the limiting reagent, if any, in the following reaction mixtures.

(i) 300 atoms of A + 200 molecules of B

(ii) 2 mol A + 3 mol B

(iii) 100 atoms of A + 100 molecules of B

(iv) 5 mol A + 2.5 mol B

(v) 2.5 mol A + 5 mol B

Limiting reagent:

It determines the extent of a reaction. It is the first to get consumed during a reaction, thus causes the reaction to stop and limits the amount of product formed.

1 atom of A reacts with 1 molecule of B. Similarly, 200 atoms of A reacts with 200 molecules of B, so 100 atoms of A are unused. Hence, B is the limiting reagent.

1 mole of A reacts with 1 mole of B. Similarly, 2 moles of A reacts with 2 moles of B, so 1 mole of B is unused. Hence, A is the limiting reagent.

(iii) 100 atoms of A + 100 molecules of Y

1 atom of A reacts with 1 molecule of Y. Similarly, 100 atoms of A reacts with 100 molecules of Y. Hence, it is a stoichiometric mixture where there is no limiting reagent.

1 mole of A reacts with 1 mole of B. Similarly 2.5 moles of A reacts with 2.5 moles of B, so 2.5 moles of A is unused. Hence, B is the limiting reagent.

(v) 2.5 mol A + 5 mol B

1 mole of A reacts with 1 mole of B. Similarly, 2.5 moles of A reacts with 2.5 moles of B, so 2.5 moles of B is unused. Hence, A is the limiting reagent.

Q24. Dinitrogen and dihydrogen react with each other to produce ammonia according to the following chemical equation: N 2 (g) + H 2 (g)→ 2NH 3 (g)

(i) Calculate the mass of \(\begin{array}{l}NH_{ 3 }\end{array} \) produced if \(\begin{array}{l}2 \; \times \;10^{ 3 }\end{array} \) g N 2 reacts with \(\begin{array}{l}1 \; \times \;10^{ 3 }\end{array} \) g of H 2 ?

 (ii) Will any of the two reactants remain unreacted?

(iii) If yes, which one and what would be its mass.

(i) 1 mol of N 2 i.e., 28 g reacts with 3 moles of H 2 i.e., 6 g of H 2

∴ 2000 g of N 2 will react with H 2 =  \(\begin{array}{l}\frac{6}{28}\times 200g = 428.6g\end{array} \)

Thus, N 2 is the limiting reagent while H 2 is the excess reagent

2 mol of N 2 i.e., 28 g of N 2 produces NH 3 = 2 mol

Therefore, 2000 g will produces NH 3 =  \(\begin{array}{l}\frac{34}{28}\times 2000 g\end{array} \)

= 2428.57 g

(ii) H 2 will remain unreacted

(iii) Mass left unreacted = 1000g – 428.6g = 571.4g

Q25. How are 0.50 mol Na 2 CO 3 and 0.50 M Na 2 CO 3 different?

Molar mass of \(\begin{array}{l}Na_{ 2 }CO_{ 3 }\end{array} \) :

= (2 × 23) + 12 + (3 × 16)

= 106 g \(\begin{array}{l}mol^{ -1 }\end{array} \)

1 mole of \(\begin{array}{l}Na_{ 2 }CO_{ 3 }\end{array} \) means 106 g of \(\begin{array}{l}Na_{ 2 }CO_{ 3 }\end{array} \)

Therefore, 0.5 mol of \(\begin{array}{l}Na_{ 2 }CO_{ 3 }\end{array} \)

= \(\begin{array}{l}\frac{ 106 \; g }{ 1 \; mol } \; \times \; 0.5 \; mol \end{array} \) \(\begin{array}{l}Na_{ 2 }CO_{ 3 }\end{array} \)

= 53 g of \(\begin{array}{l}Na_{ 2 }CO_{ 3 }\end{array} \)

0.5 M of \(\begin{array}{l}Na_{ 2 }CO_{ 3 }\end{array} \) = 0.5 mol/L \(\begin{array}{l}Na_{ 2 }CO_{ 3 }\end{array} \)

Hence, 0.5 mol of \(\begin{array}{l}Na_{ 2 }CO_{ 3 }\end{array} \) is in 1 L of water or 53 g of \(\begin{array}{l}Na_{ 2 }CO_{ 3 }\end{array} \) is in 1 L of water.

Q26. If 10 volumes of dihydrogen gas reacts with five volumes of dioxygen gas, how many volumes of water vapour would be produced?

2 volumes of dihydrogen react with 1 volume of dioxygen to produce two volumes of water vapour.

Hence, 10 volumes of dihydrogen will react with five volumes of dioxygen to produce 10 volumes of water vapour.

Q27. Convert the following into basic units:

(i) 28.7 pm

(ii) 15.15 pm

(iii) 25365 mg

1 pm = \(\begin{array}{l}10^{ -12 } \; m\end{array} \)

28.7 pm = 28.7 × \(\begin{array}{l}10^{ -12 } \; m\end{array} \)

= 2.87 × \(\begin{array}{l}10^{ -11 } \; m\end{array} \)

15.15 pm = 15.15 × \(\begin{array}{l}10^{ -12 } \; m\end{array} \)

= 1.515 × \(\begin{array}{l}10^{ -11 } \; m\end{array} \)

1 mg = \(\begin{array}{l}10^{ -3 } \; g\end{array} \)

1 mg = 10 -6 kg

25365 mg = 25365 x 10 -6 kg

25365 mg = 2.5365 × \(\begin{array}{l}10^{ -2 } \; kg\end{array} \)

Q28. Which one of the following will have the largest number of atoms?

(i) 1 g Au (s)

(ii) 1 g Na (s)

(iii) 1 g Li (s)

(iv) 1 g of \(\begin{array}{l}Cl_{ 2 }\end{array} \) (g)

(i) 1 g of Au (s)

= \(\begin{array}{l}\frac{ 1 }{ 197 }\end{array} \) mol of Au (s)

= \(\begin{array}{l}\frac{ 6.022 \; \times \; 10^{ 23 } }{ 197 }\end{array} \) atoms of Au (s)

= 3.06 \(\begin{array}{l}\times \; 10^{ 21 }\end{array} \) atoms of Au (s)

(ii) 1 g of Na (s)

= \(\begin{array}{l}\frac{ 1 }{ 23 }\end{array} \) mol of Na (s)

= \(\begin{array}{l}\frac{ 6.022 \; \times \; 10^{ 23 } }{ 23 }\end{array} \) atoms of Na (s)

= 0.262 \(\begin{array}{l}\times \; 10^{ 23 }\end{array} \) atoms of Na (s)

= 26.2 \(\begin{array}{l}\times \; 10^{ 21 }\end{array} \) atoms of Na (s)

(iii) 1 g of Li (s)

= \(\begin{array}{l}\frac{ 1 }{ 7 }\end{array} \) mol of Li (s)

= \(\begin{array}{l}\frac{ 6.022 \; \times \; 10^{ 23 } }{ 7 }\end{array} \) atoms of Li (s)

= 0.86 \(\begin{array}{l}\times \; 10^{ 23 }\end{array} \) atoms of Li (s)

= 86.0 \(\begin{array}{l}\times \; 10^{ 21 }\end{array} \) atoms of Li (s)

= \(\begin{array}{l}\frac{ 1 }{ 71 }\end{array} \) mol of \(\begin{array}{l}Cl_{ 2 }\end{array} \) (g)

(Molar mass of \(\begin{array}{l}Cl_{ 2 }\end{array} \) molecule = 35.5 × 2 = 71 g \(\begin{array}{l}mol^{ -1 }\end{array} \) )

= \(\begin{array}{l}\frac{ 6.022 \; \times \; 10^{ 23 } }{ 71 }\end{array} \) atoms of \(\begin{array}{l}Cl_{ 2 }\end{array} \) (g)

= 0.0848 \(\begin{array}{l}\times \; 10^{ 23 }\end{array} \) atoms of \(\begin{array}{l}Cl_{ 2 }\end{array} \) (g)

= 8.48 \(\begin{array}{l}\times \; 10^{ 21 }\end{array} \) atoms of \(\begin{array}{l}Cl_{ 2 }\end{array} \) (g)

Therefore, 1 g of Li (s) will have the largest no. of atoms.

Q29. Calculate the molarity of a solution of ethanol in water, in which the mole fraction of ethanol is 0.040 (assume the density of water to be one).

Mole fraction of \(\begin{array}{l}C_{ 2 }H_{ 5 }OH\end{array} \)

= \(\begin{array}{l}\frac{Number \; of \; moles \; of \; C_{ 2 }H_{ 5 }OH}{Number \; of \; moles \; of \; solution}\end{array} \)

0.040 = \(\begin{array}{l}\frac{n_{C_{ 2 }H_{ 5 }OH}}{n_{C_{ 2 }H_{ 5 }OH} \; + \; n_{H_{ 2 }O}}\end{array} \) ——(1)

No. of moles present in 1 L water:

Substituting the value of \(\begin{array}{l}n_{ H_{ 2 }O}\end{array} \) in eq (1),

0.96 \(\begin{array}{l}n_{C_{ 2 }H_{ 5 }OH}\end{array} \) = 2.222 mol

Therefore, molarity of solution

= \(\begin{array}{l}\frac{ 2.314 \; mol }{ 1 \; L }\end{array} \)

Q30. What will be the mass of one 12 C atom in g?

1 mole of carbon atoms

= \(\begin{array}{l}6.023 \; \times \; 10^{ 23 }\end{array} \) atoms of carbon

= 12 g of carbon

Therefore, mass of 1 atom of \(\begin{array}{l}_{}^{ 12 }\textrm{ C }\end{array} \)

= \(\begin{array}{l}\frac{ 12 \; g }{ 6.022 \; \times \; 10^{ 23 }}\end{array} \)

= \(\begin{array}{l}1.993 \; \times \; 10^{ -23 } g\end{array} \)

Q31. How many significant figures should be present in the answer of the following calculations?

(i) \(\begin{array}{l}\frac{ 0.02856 \; \times \; 298.15 \; \times \; 0.112}{ 0.5785 }\end{array} \)

(ii) 5 × 5.364

(iii) 0.0125 + 0.7864 + 0.0215

Least precise number = 0.112

Therefore, no. of significant numbers in the answer

= No. of significant numbers in 0.112

Least precise number = 5.364

= No. of significant numbers in 5.364

As the least no. of decimal place in each term is 4. Hence, the no. of significant numbers in the answer is also 4.

Q32. Use the data given in the following table to calculate the molar mass of naturally occurring argon isotopes:

Molar mass of Argon:

= [ \(\begin{array}{l}( 35.96755 \; \times \; \frac{ 0.337 }{ 100 })\end{array} \) + \(\begin{array}{l}( 37.96272 \; \times \; \frac{ 0.063 }{ 100 })\end{array} \) + \(\begin{array}{l}( 39.9624 \; \times \; \frac{ 99.600 }{ 100 })\end{array} \) ]

= [0.121 + 0.024 + 39.802] \(\begin{array}{l}g \; mol^{ -1 }\end{array} \)

= 39.947 \(\begin{array}{l}g \; mol^{ -1 }\end{array} \)

Q33. Calculate the number of atoms in each of the following

(i) 52 moles of Ar

(ii) 52 u of He

(iii) 52 g of He

1 mole of Ar = \(\begin{array}{l}6.023 \; \times \; 10^{ 23 }\end{array} \) atoms of Ar

Therefore, 52 moles of Ar = 52 × \(\begin{array}{l}6.023 \; \times \; 10^{ 23 }\end{array} \) atoms of Ar

= \(\begin{array}{l}3.131 \; \times \; 10^{ 25 }\end{array} \) atoms of Ar

1 atom of He = 4 u of He

4 u of He = 1 atom of He

1 u of He = \(\begin{array}{l}\frac{ 1 }{ 4 }\end{array} \) atom of He

52 u of He = \(\begin{array}{l}\frac{ 52 }{ 4 }\end{array} \) atom of He

= 13 atoms of He

4 g of He = \(\begin{array}{l}6.023 \; \times \; 10^{ 23 }\end{array} \) atoms of He

52 g of He = \(\begin{array}{l}\frac{ 6.023 \; \times \; 10^{ 23 } \; \times \;52 }{ 4 }\end{array} \) atoms of He

= \(\begin{array}{l}7.829 \; \times \; 10^{ 24 }\end{array} \) atoms of He

Q34. A welding fuel gas contains carbon and hydrogen only. Burning a small sample of it in oxygen gives 3.38 g carbon dioxide, 0.690 g of water and no other products. A volume of 10.0 L (measured at STP) of this welding gas is found to weigh 11.6 g.   Find:

(i) Empirical formula

(ii) Molar mass of the gas, and

(iii) Molecular formula

1 mole of \(\begin{array}{l}CO_{ 2 }\end{array} \) contains 12 g of carbon

Therefore, 3.38 g of \(\begin{array}{l}CO_{ 2 }\end{array} \) will contain carbon

= \(\begin{array}{l}\frac{ 12 \; g }{ 44 \; g } \; \times 3.38 \; g\end{array} \)

18 g of water contains 2 g of hydrogen

Therefore, 0.690 g of water will contain hydrogen

= \(\begin{array}{l}\frac{ 2 \; g }{ 18 \; g } \; \times 0.690\end{array} \)

As hydrogen and carbon are the only elements of the compound. Now, the total mass is:

= 0.9217 g + 0.0767 g

Therefore, % of C in the compound

= \(\begin{array}{l}\frac{ 0.9217 \; g }{ 0.9984 \; g } \; \times 100\end{array} \)

% of H in the compound

= \(\begin{array}{l}\frac{ 0.0767 \; g }{ 0.9984 \; g } \; \times 100\end{array} \)

Moles of C in the compound,

= \(\begin{array}{l}\frac{ 92.32 }{ 12.00 }\end{array} \)

Moles of H in the compound,

= \(\begin{array}{l}\frac{ 7.68 }{ 1 }\end{array} \)

Therefore, the ratio of carbon to hydrogen is,

Therefore, the empirical formula is CH.

(ii) Molar mass of the gas

Weight of 10 L of gas at STP = 11.6 g

Therefore, weight of 22.4 L of gas at STP

= \(\begin{array}{l}\frac{ 11.6 \; g }{ 10 \; L } \; \times \; 22.4 \; L\end{array} \)

Empirical formula mass:

CH = 12 + 1

n = \(\begin{array}{l}\frac{ Molar \; mass \; of \; gas}{Empirical \; formula \; mass \; of \; gas}\end{array} \)

= \(\begin{array}{l}\frac{ 26 \; g }{ 13 \; g}\end{array} \)

Therefore, molecular formula = 2 x CH = \(\begin{array}{l}C_{ 2 }H_{ 2 }\end{array} \) .

Q35. Calcium carbonate reacts with aqueous HCl to give CaCl 2 and CO 2 according to the reaction, CaCO 3 (s) + 2 HCl (aq) →  CaCl 2 (aq) + CO 2 (g) + H 2 O(l)

What mass of CaCO 3 is required to react completely with 25 mL of 0.75 M HCl?

0.75 M of HCl

≡ 0.75 mol of HCl are present in 1 L of water

≡ [(0.75 mol) × (36.5 g mol –1 )] HCl is present in 1 L of water

≡ 27.375 g of HCl is present in 1 L of water

Thus, 1000 mL of solution contains 27.375 g of HCl

Therefore, amt of HCl present in 25 mL of solution

= \(\begin{array}{l}\frac{ 27.375 \; g }{ 1000 \; mL } \; \times \; 25 \; mL\end{array} \)

Given chemical reaction,

2 mol of HCl (2 × 36.5 = 73 g) react with 1 mol of \(\begin{array}{l}CaCO_{ 3 }\end{array} \) (100 g)

Therefore, amt of \(\begin{array}{l}CaCO_{ 3 }\end{array} \) that will react with 0.6844 g

= \(\begin{array}{l}\frac{ 100 }{ 73 } \; \times \; 0.6844 \; g\end{array} \)

Q36. Chlorine is prepared in the laboratory by treating manganese dioxide (MnO 2 ) with aqueous hydrochloric acid according to the reaction:

4 HCl (aq) + MnO 2 (s) → 2H 2 O (l) + MnCl 2 (aq) + Cl 2 (g)

How many grams of HCl react with 5.0 g of manganese dioxide?

1 mole of \(\begin{array}{l}MnO_{2}\end{array} \) = 55 + 2 × 16 = 87 g

4 mole of HCl = 4 × 36.5 = 146 g

1 mole of \(\begin{array}{l}MnO_{2}\end{array} \) reacts with 4 mol of HCl

5 g of \(\begin{array}{l}MnO_{ 2 }\end{array} \) will react with:

= \(\begin{array}{l}\frac{146 \; g}{87 \; g} \; \times \; 5 \; g\end{array} \) HCl

= 8.4 g HCl

Therefore, 8.4 g of HCl will react with 5 g of \(\begin{array}{l}MnO_{2}\end{array} \) .

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Chemistry 11 Budd: Unit 7 - Stoichiometry

• Unit 1 - Intro to Chemistry
• Unit 2 - Atoms
• Unit 6 - Chemical reactions
• Unit 5 - The Mole Concept
• Unit 7 - Stoichiometry
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