stoichiometry

Chapter 4 Stoichiometry SLO based solved Exercise 9th class new syllabus

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Get comprehensive solutions for Chapter 4 Stoichiometry from the 9th Class New Syllabus. This SLO-based solved exercise is designed to help students master the concepts of stoichiometry effectively.

(i) How many atoms are present in one gram of H2O?

1 mole of H2O = 18 g and contains 6.022×1023 molecules.
1 molecule of H2O = 3 atoms (2 H + 1 O).

Thus,
Atoms in 1 gram of H2O: 6.022×1023/18×3≈10.02×1022

Correct answer: (a) 10.02×1022 atoms

(ii) Which is the correct formula of calcium phosphide?

The formula of calcium phosphide is based on the charges of Ca2+ and P3−. To balance charges, the formula is Ca3P2

Correct answer: (c) Ca3P2

(iii) How many atomic mass units (amu) are there in one gram?

1 amu = 1.66×10−24g.

Number of amu in 1 g=1/1.66×10−24≈6.022×1023

Correct answer: (b) 6.022×1023

(iv) Structural formula of 2-hexene is CH3−CH=CH−(CH2)2−CH3. What is its empirical formula?

The molecular formula of 2-hexene is C6H12
Empirical formula is obtained by dividing subscripts by their greatest common divisor (6). Empirical formula=CH2

Correct answer: (b) CH2

(v) How many moles are there in 25 g of H2SO4?

Molar mass of H2SO4 = 2+32+(4×16)=98 g/mol

Moles=Mass/Molar mass=25/98≈0.255 moles.

Correct answer: (c) 0.255 moles

(vi) A necklace has 6 g of diamonds in it. What are the number of carbon atoms in it?

Diamonds are made of carbon.
Molar mass of carbon = 12 g/mol.
Number of moles of carbon: Moles=6/12=0.5 moles

Number of carbon atoms: 0.5×6.022×1023≈3.01×1023

Correct answer: (d) 3.01×1023

(vii) What is the mass of Al in 204 g of aluminum oxide (Al2O3)?

Molar mass of Al2O3 = 2(27)+3(16)=102 g/mol
Mass fraction of Al: 2(27)/102=54/102=0.5294

Mass of Al in 204 g of Al2O3: 0.5294×204≈108 g

Correct answer: (d) 108 g

(viii) Which one of the following compounds will have the highest percentage by mass of nitrogen?

  • For CO(NH2)2(urea):
    Molar mass = 12+16+2(14)+4(1)=60
    Nitrogen mass = 28
    Percentage of N = 28/60×100=46.67%
  • For NH4NO3:
    Molar mass = 14+4+14+3(16)=80
    Nitrogen mass = 28.
    Percentage of N = 28/80×100=35%

Correct answer: (a) CO(NH2)2

(ix) When one mole of each of the following compounds is reacted with oxygen, which will produce the maximum amount of CO2?

The number of moles of CO2 depends on the number of carbon atoms in the compound. Among the given options:

  • Methane (CH4): Produces 1 mole of CO2.
  • Ethane (C2H6): Produces 2 moles of CO2.
  • Diamond (C): Produces 1 mole of CO2

Correct answer: (b) Ethane

(x) What mass of 95% Ca(OH)2 will be required to neutralize 50 cm3 of 0.5 M H2SO4?

  • Reaction: Ca(OH)2+H2SO4→CaSO4+2H2O
    1 mole of H2SO4 reacts with 1 mole of Ca(OH)2.
    Moles of H2SO4: 0.5×0.05=0.025 moles
    Mass of pure Ca(OH)2: 0.025×(40+2(16+1))=0.025×74=1.85 g.
  • Mass of 95% pure Ca(OH)2: 1.850.95≈1.95 g

Correct answer: (d) 1.95 g

2. Questions for Short Answers

i. Write down the chemical formula of barium nitride.
The chemical formula of barium nitride is Ba₃N₂.
Explanation: Barium (Ba) is an alkaline earth metal with a +2 oxidation state, and nitride (N) is a nonmetal with a -3 charge. To balance the charges, three barium ions combine with two nitride ions, forming Ba₃N₂.

ii. Find out the molecular formula of a compound whose empirical formula is CH₂O and its molar mass is 180.
Step 1: Calculate the molar mass of the empirical formula (CH₂O).
Molar mass of CH₂O = 12 (C) + 2(1) (H) + 16 (O) = 30 g/mol.

Step 2: Determine the ratio of the molecular formula to the empirical formula.
Ratio=Molar mass of compound / Molar mass of empirical formula=180/30=6.

Step 3: Multiply the empirical formula by the ratio.
Molecular formula = (CH2O)6=C6H12O6 = C₆H₁₂O₆.
The molecular formula of the compound is C₆H₁₂O₆ (glucose).

iii. How many molecules are present in 1.5 g H₂O?
Step 1: Calculate the molar mass of water (H₂O).
Molar mass = 2(1) + 16 = 18 g/mol.

Step 2: Find the number of moles in 1.5 g of H₂O.
Moles = Mass/Molar mass=1.5/18=0.0833 mol

Step 3: Calculate the number of molecules using Avogadro’s number (6.022×1023 molecules/mol)
Number of molecules = 0.0833×6.022×1023=5.02×1022
Thus, there are approximately 5.02×1022

iv. What is the difference between a mole and Avogadro’s number?

  • Mole: A mole is a unit in chemistry that represents a quantity of substance. One mole of any substance contains the same number of particles (atoms, molecules, or ions) as there are in 12 g of carbon-12.
  • Avogadro’s Number: Avogadro’s number is a constant that defines the number of particles in one mole of a substance. It is equal to 6.022×1023 particles/mol

Difference: The mole is a quantity (like a dozen), while Avogadro’s number provides the specific count of particles in one mole.

v. Write down the chemical equation for the given reaction.
Copper + Sulphuric acid → Copper sulphate + Sulphur dioxide + Water
Balanced chemical equation:
Cu+2H₂SO₄→CuSO₄+SO₂+2H₂O

3. Constructed Response Questions

i. Different compounds will never have the same molecular formula but can have the same empirical formula. Explain.

  • The empirical formula represents the simplest whole-number ratio of elements in a compound, whereas the molecular formula shows the exact number of atoms of each element in a molecule.
  • Example: Acetic acid (C₂H₄O₂) and glucose (C₆H₁₂O₆) have the same empirical formula (CH₂O), but their molecular formulas are different.
  • Explanation: Compounds with different molecular structures can share the same empirical formula because the molecular formula is a multiple of the empirical formula.

ii. Write down the chemical formulas of the following compounds.

  1. Calcium phosphate: Ca₃(PO₄)₂
  2. Aluminium nitride: AlN
  3. Sodium acetate: CH₃COONa
  4. Ammonium carbonate: (NH₄)₂CO₃
  5. Bismuth sulphate: Bi₂(SO₄)₃

iii. Why does Avogadro’s number have immense importance in chemistry?
Avogadro’s number (6.022×1023 is essential in chemistry because:

  • It establishes a connection between the microscopic (atoms/molecules) and macroscopic (grams/moles) scales.
  • It allows chemists to calculate the number of particles in a given mass of substance.
  • It is crucial for determining stoichiometric relationships in chemical reactions, enabling precise calculations for reactants and products.

iv. When 8.657 g of a compound were converted into elements, it gave 5.217 g of carbon, 0.962 g of hydrogen, and 2.478 g of oxygen. Calculate the percentage of each element present in this compound.
Step 1: Calculate the total mass of the compound.
Total mass = 5.217+0.962+2.478=8.657 g

Step 2: Calculate the percentage of each element.

  • Carbon: 5.2178.657×100=60.3%
  • Hydrogen: 0.9628.657×100=11.1%
  • Oxygen: 2.4788.657×100=28.6%

Thus, the percentages are:

  • Carbon: 60.3%
  • Hydrogen: 11.1%
  • Oxygen: 28.6%.

4. Descriptive Questions

i. Which conditions must be fulfilled before writing a chemical equation for a reaction?
To write a chemical equation, the following conditions must be fulfilled:

  1. Correct identification of reactants and products: Ensure the substances involved in the reaction are accurately identified.
  2. Balancing the equation: The equation must follow the law of conservation of mass, where the number of atoms for each element is the same on both sides.
  3. Physical states of substances: Indicate the physical states of reactants and products (solid (s), liquid (l), gas (g), or aqueous (aq)).
  4. Reaction conditions: Specify the conditions such as temperature, pressure, or the use of catalysts if required for the reaction.

ii. Explain the concepts of Avogadro’s number and mole.

  • Avogadro’s Number: Avogadro’s number (6.022×1023 is the number of particles (atoms, molecules, or ions) in one mole of a substance. It bridges the gap between the microscopic scale and the macroscopic scale.
  • Mole: A mole is a unit of measurement used in chemistry to represent 6.022×1023 particles of a substance. It is the amount of a substance that contains as many entities as there are atoms in 12 grams of carbon-12.

iii. How many grams of CO₂ will be produced when we react 10 g of CH₄ with excess of O₂ according to the following equation?
CH4+2O2→CO2+2H2OCH₄ + 2O₂ → CO₂ + 2H₂O

  1. Molar masses:
    • CH₄ = 16 g/mol
    • CO₂ = 44 g/mol
  2. Moles of CH₄:
    Moles = Mass/Molar mass=10/16=0.625 mol
  3. Moles of CO₂:
    From the equation, 1 mole of CH₄ produces 1 mole of CO₂.
    Moles of CO₂ = 0.625 mol.
  4. Mass of CO₂:
    Mass = Moles × Molar mass = 0.625×44=27.5 g

Answer: 27.5 g of CO₂ will be produced.

iv. How many moles of coal are needed to produce 10 moles of CO according to the following equation?
3C+O2+H2O→H2+3CO3C + O₂ + H₂O → H₂ + 3CO

From the equation, 3 moles of carbon (C) produce 3 moles of CO.
Thus, to produce 10 moles of CO: Moles of coal=103=3.33 moles.\text{Moles of coal} = \frac{10}{3} = 3.33 \, \text{moles}.

Answer: 3.33 moles of coal are needed.

v. How much SO₂ is needed in grams to produce 10 moles of sulphur?
2H2S+SO2→2H2O+3S

From the equation, 1 mole of SO₂ produces 1.5 moles of sulphur.
To produce 10 moles of sulphur: Moles of SO₂=10/1.5=6.67 moles

Molar mass of SO₂ = 32 (S) + 2(16) = 64 g/mol.
Mass = Moles × Molar mass = 6.67×64=426.88 g

Answer: 426.88 g of SO₂ are needed.

vi. How much ammonia is needed in grams to produce 1 kg of urea fertilizer?
2NH3+CO2→CO(NH2)2+H2O

  1. Molar masses:
    • NH₃ = 17 g/mol
    • CO(NH₂)₂ (urea) = 60 g/mol
  2. Moles of urea:
    Moles = Mass/Molar mass=1000/60=16.67 mol
  3. Moles of NH₃:
    From the equation, 2 moles of NH₃ produce 1 mole of urea.
    Moles of NH₃ = 16.67×2=33.34 mol.
  4. Mass of NH₃:
    Mass = Moles × Molar mass = 33.34×17=566.78 g

Answer: 566.78 g of ammonia is needed.

vii. Calculate the number of atoms in the following:

(a) 3 g of H₂:
Moles of H₂ = Mass/Molar mass=3/2=1.5 mol
Number of molecules = 1.5×6.022×1023=9.033×1023
Each molecule of H₂ contains 2 atoms, so:
Number of atoms = 9.033×1023×2=1.8066×1024
Number of molecules = 3.4×6.022×1023=2.0475×1024
Each molecule of N₂ contains 2 atoms, so:
Number of atoms = 2.0475×1024×2=4.095×1024
Molar mass of C₆H₁₂O₆ = 6(12)+12(1)+6(16)=180 g/mol
Number of molecules = 0.0556×6.022×1023=3.348×1022
Each molecule contains 6+12+6=246 + 12 + 6 = 24 atoms.
Number of atoms = 3.348×1022×24=8.035×1023

5. Investigative Questions

i. How many moles of water are needed for a single adult?
Volume of water per glass = 400 cm3
Density of water = 1 g/cm3
Mass of water = 400×8=3200 g

Molar mass of water = 18 g/mol.
Moles of water = Mass/Molar mass=3200/18=177.78 mol

Answer: 177.78 moles of water are needed.

ii. How many moles of SiO₂ are present in the glass?
Assuming the mass of glass is provided, moles = Mass of glass/Molar mass of SiO₂

where Molar mass of SiO₂=28+2(16)=60 g/mol

Stoichiometry Federal board mcqs

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Stoichiometry

Practice Stoichiometry MCQs designed for Federal Board students. Enhance your problem-solving skills with topic-specific questions and detailed explanations.

Mole, molar volume, molar mass, and the density of gases

  1. What is defined as one mole of any substance?
  • A) 6.02 x 10²³ molecules
  • B) The mass of an atom
  • C) One liter of a gas at STP
  • D) One gram of a substance
  • Answer: A) 6.02 x 10²³ molecules
  1. The molar mass of a substance is:
  • A) The volume of one mole of gas at STP
  • B) The mass of one mole of a substance
  • C) The number of moles per liter
  • D) The density of the substance
  • Answer: B) The mass of one mole of a substance
  1. What is the molar volume of an ideal gas at STP?
  • A) 22.4 L
  • B) 6.02 L
  • C) 1.00 L
  • D) 44.8 L
  • Answer: A) 22.4 L
  1. Which statement best defines Avogadro’s number?
  • A) The volume occupied by 1 mole of any gas at STP
  • B) The number of atoms in 12 g of carbon-12
  • C) The atomic mass of any element
  • D) The molar mass of a substance
  • Answer: B) The number of atoms in 12 g of carbon-12
  1. The density of a gas at STP can be calculated using:
  • A) Density = molar mass / molar volume
  • B) Density = molar mass x volume
  • C) Density = moles x temperature
  • D) Density = molar mass / pressure
  • Answer: A) Density = molar mass / molar volume
  1. If a gas has a molar mass of 44 g/mol, its density at STP would be:
  • A) 0.5 g/L
  • B) 1.96 g/L
  • C) 2.5 g/L
  • D) 44 g/L
  • Answer: B) 1.96 g/L
  1. Which gas law explains the relationship between molar volume and temperature?
  • A) Boyle’s Law
  • B) Charles’s Law
  • C) Avogadro’s Law
  • D) Ideal Gas Law
  • Answer: B) Charles’s Law
  1. At STP, which of the following gases will have the greatest density?
  • A) ( \text{H}_2 )
  • B) ( \text{O}_2 )
  • C) ( \text{CO}_2 )
  • D) ( \text{He} )
  • Answer: C) ( \text{CO}_2 )
  1. One mole of ( \text{H}_2 ) gas at STP occupies:
  • A) 11.2 L
  • B) 22.4 L
  • C) 33.6 L
  • D) 44.8 L
  • Answer: B) 22.4 L
  1. The mass of 1 mole of ( \text{CO}_2 ) is approximately:
    • A) 28 g
    • B) 32 g
    • C) 44 g
    • D) 16 g
    • Answer: C) 44 g
  2. Which term describes the mass of one mole of any chemical element or compound?
    • A) Molecular weight
    • B) Molar volume
    • C) Atomic mass
    • D) Molar mass
    • Answer: D) Molar mass
  3. If a gas has a molar mass of 2 g/mol, what would be its density at STP?
    • A) 0.089 g/L
    • B) 0.178 g/L
    • C) 0.5 g/L
    • D) 2 g/L
    • Answer: B) 0.178 g/L
  4. The molar mass of ( \text{O}_2 ) is:
    • A) 16 g/mol
    • B) 32 g/mol
    • C) 8 g/mol
    • D) 48 g/mol
    • Answer: B) 32 g/mol
  5. At constant temperature and pressure, the volume of gas is directly proportional to:
    • A) Mass
    • B) Molar mass
    • C) Number of moles
    • D) Density
    • Answer: C) Number of moles
  6. What is the volume occupied by 0.5 moles of a gas at STP?
    • A) 11.2 L
    • B) 22.4 L
    • C) 5.6 L
    • D) 44.8 L
    • Answer: A) 11.2 L
  7. The density of a gas at STP depends on:
    • A) Temperature
    • B) Molar mass
    • C) Volume
    • D) Avogadro’s number
    • Answer: B) Molar mass
  8. Which unit is used to measure molar volume of a gas?
    • A) g/mol
    • B) mol/L
    • C) L/mol
    • D) mol/g
    • Answer: C) L/mol
  9. For a given gas, doubling the pressure at constant temperature will:
    • A) Double the volume
    • B) Halve the volume
    • C) Keep the volume constant
    • D) Double the density
    • Answer: B) Halve the volume
  10. Which of the following represents the correct molar volume of a gas at STP?
    • A) 11.2 L/mol
    • B) 22.4 L/mol
    • C) 33.6 L/mol
    • D) 44.8 L/mol
    • Answer: B) 22.4 L/mol
  11. How many moles are present in 44.8 L of ( \text{CO}_2 ) gas at STP?
    • A) 1 mole
    • B) 2 moles
    • C) 0.5 moles
    • D) 3 moles
    • Answer: B) 2 moles
  12. If the molar mass of a gas is known, its density at STP can be calculated by:
    • A) Dividing molar mass by Avogadro’s number
    • B) Dividing molar mass by molar volume
    • C) Multiplying molar mass by volume
    • D) Dividing molar mass by temperature
    • Answer: B) Dividing molar mass by molar volume
  13. The mass of 1 mole of ( \text{N}_2 ) gas is:
    • A) 14 g
    • B) 28 g
    • C) 2 g
    • D) 32 g
    • Answer: B) 28 g
  14. Which of the following represents Avogadro’s law?
    • A) ( P \propto V )
    • B) ( V \propto T )
    • C) ( V \propto n )
    • D) ( P \propto T )
    • Answer: C) ( V \propto n )
  15. The density of ( \text{O}_2 ) gas at STP is approximately:
    • A) 1.43 g/L
    • B) 0.5 g/L
    • C) 2.86 g/L
    • D) 1.2 g/L
    • Answer: A) 1.43 g/L
  16. What volume will 1 mole of ( \text{CH}_4 ) gas occupy at STP?
    • A) 44.8 L
    • B) 22.4 L
    • C) 11.2 L
    • D) 5.6 L
    • Answer: B) 22.4 L

Stoichiometric calculations, mole ratios, and mole-mole calculations

  1. In a balanced chemical equation, the coefficients represent the:
  • A) Mass ratio of reactants and products
  • B) Mole ratio of reactants and products
  • C) Volume ratio at STP
  • D) Molecular weights
  • Answer: B) Mole ratio of reactants and products
  1. What is the mole ratio of H2 to O2 in the reaction 2H2 + O2 → 2H2O?
  • A) 1:1
  • B) 2:1
  • C) 1:2
  • D) 3:2
  • Answer: B) 2:1
  1. In the equation N2 + 3H2 → 2NH3 , how many moles of NH3 are produced from 6 moles of H2 ?
  • A) 1 mole
  • B) 2 moles
  • C) 3 moles
  • D) 4 moles
  • Answer: C) 3 moles
  1. What is the first step in solving stoichiometric problems?
  • A) Calculating the molar mass
  • B) Balancing the chemical equation
  • C) Converting grams to moles
  • D) Determining the limiting reagent
  • Answer: B) Balancing the chemical equation
  1. In a chemical reaction, the mole ratio between two reactants can be determined from:
  • A) Their atomic masses
  • B) The coefficients in the balanced equation
  • C) Their physical states
  • D) The reaction rate
  • Answer: B) The coefficients in the balanced equation
  1. Using the balanced equation 4Fe + 3O2 → 2Fe2O3 , how many moles of Fe2O3 ) will form from 4 moles of Fe?
  • A) 1 mole
  • B) 2 moles
  • C) 3 moles
  • D) 4 moles
  • Answer: B) 2 moles
  1. In the reaction ( 2KClO3 → 2KCl + 3O2 ), how many moles of O2 ) are produced from 4 moles of KClO3 )?
  • A) 2 moles
  • B) 3 moles
  • C) 4 moles
  • D) 6 moles
  • Answer: D) 6 moles
  1. If 2 moles of H2 react with O2 to form water, what is the mole ratio of H2 to H2O?
  • A) 1:1
  • B) 1:2
  • C) 2:2
  • D) 2:1
  • Answer: A) 1:1
  1. In a reaction, if the mole ratio between reactant A and product B is 1:3, then 2 moles of A will produce how many moles of B?
  • A) 1 mole
  • B) 2 moles
  • C) 3 moles
  • D) 6 moles
  • Answer: D) 6 moles
  1. The stoichiometric coefficient indicates:
    • A) The mass of each reactant and product
    • B) The energy released in the reaction
    • C) The number of moles involved in the reaction
    • D) The color of the compounds
    • Answer: C) The number of moles involved in the reaction
  2. In a combustion reaction, if the mole ratio of CH4 to O2 is 1:2, how many moles of O2 are required to completely burn 5 moles of CH4 ?
    • A) 2.5 moles
    • B) 5 moles
    • C) 10 moles
    • D) 15 moles
    • Answer: C) 10 moles
  3. Using the equation 2SO2 + O2 → 2SO3, if 2 moles of SO2 react, how many moles of SO3 will be formed?
    • A) 1 mole
    • B) 2 moles
    • C) 3 moles
    • D) 4 moles
    • Answer: B) 2 moles
  4. How many moles of HCl are needed to react completely with 1 mole of NaOH in the reaction HCl + NaOH → NaCl + H2O?
    • A) 0.5 moles
    • B) 1 mole
    • C) 1.5 moles
    • D) 2 moles
    • Answer: B) 1 mole
  5. In the reaction ( 3H2 + N2 → 2NH3, the mole ratio of N2 to NH3 is:
    • A) 1:1
    • B) 1:2
    • C) 2:3
    • D) 3:1
    • Answer: B) 1:2
  6. What is the mole ratio of CO2 ) to C2H4 ) in the complete combustion reaction C2H4 + 3O2→ 2CO2 + 2H2O?
    • A) 1:2
    • B) 2:1
    • C) 1:1
    • D) 3:2
    • Answer: B) 2:1
  7. The mole ratio of O2 + CO2 in the reaction CH4 + 2O2 → CO2 + 2H2O is:
    • A) 1:1
    • B) 1:2
    • C) 2:1
    • D) 2:3
    • Answer: A) 1:1
  8. In the balanced equation C6H12O6 → 2C2H5OH + 2CO2 , how many moles of CO2 are produced from 1 mole of C6H12O6 ?
    • A) 1 mole
    • B) 2 moles
    • C) 3 moles
    • D) 4 moles
    • Answer: B) 2 moles
  9. The equation 2Mg + O2 → 2MgO indicates that 4 moles of Mg will react with how many moles of O2 ?
    • A) 1 mole
    • B) 2 moles
    • C) 3 moles
    • D) 4 moles
    • Answer: B) 2 mole

  1. In the reaction H2 + Cl2 → 2HCl, the mole ratio of H2 to HCl is:
    • A) 1:1
    • B) 1:2
    • C) 2:1
    • D) 2:2
    • Answer: B) 1:2

Solution Stoichiometry

  1. What is solution stoichiometry used for?
  • A) To find gas pressure
  • B) To calculate solute amount in a solution
  • C) To determine atomic mass
  • D) To measure density
  • Answer: B) To calculate solute amount in a solution
  1. In solution stoichiometry, molarity (M) is defined as:
  • A) Moles of solute per liter of solution
  • B) Moles of solvent per liter of solution
  • C) Grams per liter
  • D) Moles per kilogram
  • Answer: A) Moles of solute per liter of solution
  1. If 2 moles of NaCl are dissolved in 1 L of water, what is the molarity?
  • A) 1 M
  • B) 2 M
  • C) 0.5 M
  • D) 4 M
  • Answer: B) 2 M
  1. The unit of molarity is:
  • A) mol/g
  • B) mol/L
  • C) L/mol
  • D) g/mol
  • Answer: B) mol/L
  1. To dilute a solution, you would:
  • A) Add more solute
  • B) Increase temperature
  • C) Add more solvent
  • D) Remove solvent
  • Answer: C) Add more solvent
  1. In the equation ( \text{M}_1\text{V}_1 = \text{M}_2\text{V}_2 ), ( \text{M}_2 ) represents:
  • A) Final molarity
  • B) Initial volume
  • C) Final mass
  • D) Initial molarity
  • Answer: A) Final molarity
  1. What is the molarity of a solution with 0.5 moles of KCl in 0.25 L?
  • A) 1 M
  • B) 2 M
  • C) 0.25 M
  • D) 0.5 M
  • Answer: B) 2 M
  1. If you mix equal volumes of 1 M HCl and 1 M NaOH, the resulting solution is:
  • A) Acidic
  • B) Neutral
  • C) Basic
  • D) Salty
  • Answer: B) Neutral
  1. To prepare 1 L of 1 M NaOH, you need:
  • A) 1 g NaOH
  • B) 10 g NaOH
  • C) 40 g NaOH
  • D) 0.1 g NaOH
  • Answer: C) 40 g NaOH
  1. If you dilute 1 L of 2 M solution to 2 L, the new molarity is:
    • A) 4 M
    • B) 2 M
    • C) 1 M
    • D) 0.5 M
    • Answer: C) 1 M
  2. The molarity of a solution with 3 moles of solute in 3 L is:
    • A) 0.5 M
    • B) 1 M
    • C) 2 M
    • D) 3 M
    • Answer: B) 1 M
  3. What volume of 0.5 M H2SO4 is needed for 0.25 moles of H2SO4 ?
    • A) 0.5 L
    • B) 0.25 L
    • C) 1 L
    • D) 0.75 L
    • Answer: A) 0.5 L
  4. If 0.2 L of 1 M NaCl is diluted to 1 L, the molarity becomes:
    • A) 0.2 M
    • B) 0.5 M
    • C) 1 M
    • D) 0.1 M
    • Answer: D) 0.1 M
  5. How many moles are in 250 mL of 2 M HCl?
    • A) 0.25 mol
    • B) 0.5 mol
    • C) 1 mol
    • D) 2 mol
    • Answer: B) 0.5 mol
  6. The equation M1V1 = M2V2 is used for:
    • A) Dilution calculations
    • B) Gas laws
    • C) Solid solubility
    • D) Stoichiometry only
    • Answer: A) Dilution calculations
  7. How many liters of 1 M solution contain 2 moles of solute?
    • A) 0.5 L
    • B) 1 L
    • C) 2 L
    • D) 3 L
    • Answer: C) 2 L
  8. In solution stoichiometry, molarity is used to calculate:
    • A) Mass of a gas
    • B) Volume of solution
    • C) Concentration of solution
    • D) Density
    • Answer: C) Concentration of solution
  9. How many moles are present in 0.5 L of a 1 M solution?
    • A) 1 mole
    • B) 0.5 mole
    • C) 0.25 mole
    • D) 2 moles
    • Answer: B) 0.5 mole
  10. To dilute a 2 M solution to 1 M, you need to:
    • A) Add twice the amount of solvent
    • B) Remove half the solute
    • C) Add solute
    • D) Heat the solution
    • Answer: A) Add twice the amount of solvent
  11. A 1 M solution of HCl contains:
    • A) 1 mole of HCl in 1 L of solution
    • B) 1 gram of HCl in 1 L
    • C) 2 moles of HCl
    • D) 1 mole of HCl in 100 mL
    • Answer: A) 1 mole of HCl in 1 L of solution

Limiting and non-limiting reactants:

  1. What is a limiting reactant?
  • A) The reactant that is completely used up
  • B) The reactant in excess
  • C) The product formed
  • D) The reactant that remains
  • Answer: A) The reactant that is completely used up
  1. What is a non-limiting reactant?
  • A) Reactant completely used up
  • B) Reactant that is left over
  • C) Reactant that forms no product
  • D) Product formed
  • Answer: B) Reactant that is left over
  1. In a reaction, the limiting reactant determines:
  • A) Only the products
  • B) The amount of product formed
  • C) Reaction rate
  • D) The final color of the solution
  • Answer: B) The amount of product formed
  1. If all of a reactant is used up, it is:
  • A) Excess reactant
  • B) Limiting reactant
  • C) Product
  • D) Catalyst
  • Answer: B) Limiting reactant
  1. Which reactant is in excess?
  • A) Reactant fully consumed
  • B) Reactant partially left over
  • C) Reactant that forms the most product
  • D) Product formed in a reaction
  • Answer: B) Reactant partially left over
  1. How is the limiting reactant identified?
  • A) By mass alone
  • B) By comparing moles needed vs. moles available
  • C) By color of solution
  • D) By boiling point
  • Answer: B) By comparing moles needed vs. moles available
  1. What happens to the excess reactant?
  • A) It forms all products
  • B) It remains after the reaction
  • C) It becomes limiting
  • D) It is used first
  • Answer: B) It remains after the reaction
  1. Why is the limiting reactant important?
  • A) Determines reaction color
  • B) Limits amount of product
  • C) Increases temperature
  • D) Changes reaction rate
  • Answer: B) Limits amount of product
  1. In the reaction 2H2 + O2 → 2H2O, if (O2) is limited, which is in excess?
  • A) H2
  • B) H2O
  • C) O2
  • D) All are limiting
  • Answer: A) H2
  1. To find the limiting reactant, you must know:
    • A) Molar masses of products
    • B) Mole ratios from the balanced equation
    • C) Initial pressure of gases
    • D) Density of products
    • Answer: B) Mole ratios from the balanced equation

Theoretical yield, actual yield, and percent yield:

  1. What is the theoretical yield?
  • A) Actual product formed
  • B) Maximum product possible
  • C) Excess reactant remaining
  • D) Minimum product possible
  • Answer: B) Maximum product possible
  1. What is the actual yield?
  • A) Maximum yield calculated
  • B) Yield obtained in the lab
  • C) Theoretical prediction
  • D) Half of the expected yield
  • Answer: B) Yield obtained in the lab
  1. Percent yield is calculated as:
  • A) (Theoretical yield / Actual yield) x 100
  • B) (Actual yield / Theoretical yield) x 100
  • C) Actual yield + Theoretical yield
  • D) (Theoretical yield – Actual yield) x 100
  • Answer: B) (Actual yield / Theoretical yield) x 100
  1. If the actual yield is equal to the theoretical yield, percent yield is:
  • A) 50%
  • B) 100%
  • C) 0%
  • D) 200%
  • Answer: B) 100%
  1. A reaction’s theoretical yield depends on:
  • A) Limiting reactant
  • B) Excess reactant
  • C) Density
  • D) Product color
  • Answer: A) Limiting reactant
  1. Which yield is usually lower due to loss in experiments?
  • A) Theoretical yield
  • B) Actual yield
  • C) Percent yield
  • D) Limiting yield
  • Answer: B) Actual yield
  1. If theoretical yield is 10g and actual yield is 7g, percent yield is:
  • A) 70%
  • B) 140%
  • C) 30%
  • D) 10%
  • Answer: A) 70%
  1. Percent yield above 100% indicates:
  • A) No product was formed
  • B) Calculation error or impurities
  • C) Perfect reaction efficiency
  • D) Theoretical yield exceeded
  • Answer: B) Calculation error or impurities
  1. A low percent yield could indicate:
  • A) Very high efficiency
  • B) Loss of product during reaction
  • C) Overestimation of reactants
  • D) Increase in reactant mass
  • Answer: B) Loss of product during reaction
  1. Percent yield is useful to:
    • A) Measure reactant purity
    • B) Assess reaction efficiency
    • C) Calculate reaction color
    • D) Determine reaction temperature
    • Answer: B) Assess reaction efficiency

Importance of stoichiometry in the production and dosage of medicines:

  1. Stoichiometry helps in calculating:
  • A) Dosage accuracy
  • B) Medicine color
  • C) Taste of drugs
  • D) Shelf life
  • Answer: A) Dosage accuracy
  1. Why is stoichiometry important in medicine production?
  • A) To increase weight
  • B) For precise formulation
  • C) To reduce side effects
  • D) For color consistency
  • Answer: B) For precise formulation
  1. In drug formulation, stoichiometry ensures:
  • A) Stability of medicine
  • B) Accurate active ingredient amount
  • C) Reduced cost
  • D) Faster production
  • Answer: B) Accurate active ingredient amount
  1. Incorrect stoichiometric calculations can lead to:
  • A) Reduced potency
  • B) Exact dosage
  • C) Increased potency
  • D) Only taste change
  • Answer: A) Reduced potency
  1. Which concept is essential for determining the correct medicine dose?
  • A) Stoichiometry
  • B) Color theory
  • C) Surface tension
  • D) Filtration
  • Answer: A) Stoichiometry
  1. Stoichiometry helps in avoiding:
  • A) Medicine overdose
  • B) Improved packaging
  • C) Medicine odor
  • D) None of the above
  • Answer: A) Medicine overdose
  1. In pharmaceutical production, stoichiometry ensures:
  • A) Standardized drug potency
  • B) Different potency in each batch
  • C) Faster reaction times
  • D) Reduced reactivity
  • Answer: A) Standardized drug potency
  1. Accurate stoichiometric calculations in medicine are crucial for:
  • A) Patient safety
  • B) Reducing cost
  • C) Enhancing color
  • D) Speeding up reactions
  • Answer: A) Patient safety
  1. Using stoichiometry in medicine dosage prevents:
  • A) Side effects from wrong dosages
  • B) Faster dissolution
  • C) Taste change
  • D) Medicine discoloration
  • Answer: A) Side effects from wrong dosages
  1. Proper stoichiometric ratios in drugs affect:
    • A) Effectiveness and safety
    • B) Flavor of medicine
    • C) Cost of production
    • D) Packaging quality
    • Answer: A) Effectiveness and safety

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