BlogChemistryHSC Chemistry Module 2: Introduction to Quantitative Chemistry Practice Questions
HSC Chemistry Module 2: Introduction to Quantitative Chemistry Practice Questions
Struggling to find relevant and useful questions to ace your Year 11 Chemistry Module 2: Introduction to Quantitative Chemistry exam?
Look no further!
Useful practice questions like these, and the neat tips and solutions I will provide, helped me ace the HSC and boosted my student’s scores to that Band 6 range.
Now, I’ve put together my top 20 questions from Module 2 to help YOU do the same.
With varying difficulty levels (watch out for the ⭐️’s!), you can start at your current level and work your way up.
Don’t wait until it’s too late to master this module! Let our experienced HSC Chemistry Tutorssupport your studies at our Campuses in Hornsby, Chatswood or the Hills, at your home or online.
Following gravimetric analysis, you have found that your experimental mass value is lower than the theoretical mass. Account for this difference with two explanations (3 marks).
(L1.1: Conduct practical investigations to observe and measure the quantitative relationships of chemical reactions, including but not limited to masses of solids and/or liquids in chemical reactionandvolumes of gases in chemical reactions)
Question 2 ⭐⭐
If one of the products of a chemical reaction is a gas, the total volume that the atoms occupy increases after a reaction. Is this true? Explain why or why not. (3 marks)
(L1.1: Conduct practical investigations to observe and measure the quantitative relationships of chemical reactions, including but not limited to masses of solids and/or liquids in chemical reactionandvolumes of gases in chemical reactions)
Question 3 ⭐⭐⭐
Explain how you would investigate the Law of Conservation through an experiment that involves quantifying the volumes of gases in chemical reactions. (5 marks)
(L1.1: Conduct practical investigations to observe and measure the quantitative relationships of chemical reactions, including but not limited to masses of solids and/or liquids in chemical reactionandvolumes of gases in chemical reactions)
Question 4 ⭐⭐
Balance the following equations. (4 marks)
(L1.2: Relate stoichiometry to the law of conservation of mass in chemical reactions by investigating balancing chemical equations (ACSCH039) and solving problems regarding mass changes in chemical reactions (ACSCH046))
Question 5 ⭐⭐
A student combines 154 grams of carbon tetrachloride and an unknown quantity of bromine in a sealed container to produce 243 grams of dibromodichlormethane and 71 grams of chlorine. How much chlorine was used in the reaction, assuming the reactants are completely used up? (3 marks)
Question sourced from sciencing
(L1.2: Relate stoichiometry to the law of conservation of mass in chemical reactions by investigating balancing chemical equations and solving problems regarding mass changes in chemical reactions)
Mole Concept
Question 6 ⭐⭐⭐🔥
a) Design an experiment that aims to obtain the molar mass of magnesium. (6 marks)
b) Using the following experimental data, calculate the experimental molar mass of magnesium. Assume SLC.
Experimental data: m(Mg) = 0.22g, V(H2) = 240mL
(L2.1: conduct a practical investigation to demonstrate and calculate the molar mass (mass of one mole) of an element, a compound)
Question 7 ⭐⭐
A sample of 4.2g CO2gas was found to occupy 2.63L of volume. Explain how the “Ideal Gas Law” can be used to obtain the molar mass of CO2. Calculate the molar mass of CO2. (4 marks)
(L2.1: conduct a practical investigation to demonstrate and calculate the molar mass (mass of one mole) of an element, a compound)
Question 8 ⭐⭐⭐🔥
Design an experiment that tests whether the following reactions carry on in simple whole number ratios. (4 marks)
(L2.2: conduct an investigation to determine that chemicals react in simple whole number ratios by moles)
Question 9 ⭐
Answer the following questions: (6 marks)
(L2.3: explore the concept of the mole and relate this to Avogadro’s constant to describe, calculate and manipulate masses, chemical amounts and numbers of particles in: moles of elements and compounds ? = ?/Mm (n = chemical amount in moles, m = mass in grams, ?/m= molar mass in gmol-1)
Question 10 ⭐⭐
Polymers are large molecules composed of simple units repeated many times. Thus, they often have relatively simple empirical formulas. Calculate the empirical formulas of the following polymers. (5 marks)
a) Lucite (Plexiglas); 59.9% C, 8.06% H, 32.0% O
b) Saran; 24.8% C, 2.0% H, 73.1% Cl
c) polyethylene; 86% C, 14% H
d) polystyrene; 92.3% C, 7.7% H
e) Orlon; 67.9% C, 5.70% H, 26.4% N
Question sourced from Chemistry Libretexts
(L2.3: Explore the concept of the mole and relate this to Avogadro’s constant to describe, calculate and manipulate masses, chemical amounts and numbers of particles in percentage composition calculations and empirical formulae)
Question 11 ⭐⭐⭐🔥
Souring of wine occurs when ethanol is converted to acetic acid by oxygen by the following reaction:
A 1.00 L bottle of wine, labeled as 8.5% (by volume) ethanol, is found to have a defective seal. Analysis of 1.00 mL showed that there were 0.0274 grams of acetic acid in that 1.00 mL. The density of ethanol is 0.816 g/mL and the density of water is 1.00 g/mL.
a) What mass of oxygen must have leaked into the bottle? (3 marks)
b) What is the percent yield for the conversion of ethanol to acetic acid if O2 is in excess? (4 marks)
Question sourced from Washington University in St. Louis Department of Chemistry
(L2.3: explore the concept of the mole and relate this to Avogadro’s constant to describe, calculate and manipulate masses, chemical amounts and numbers of particles in limiting reagent reactions)
Concentration and Molarity
Question 12 ⭐
Explain how you have used gravimetric analysis to find the concentration of a solution of NaOH. (3 marks)
(L3.1: conduct practical investigations to determine the concentrations of solutions and investigate the different ways in which concentrations are measured)
Question 13 ⭐⭐
During gravimetric analysis, the experimental value was found to be less than the theoretical concentration of the solution. Explain THREE sources of error and suggest ways to prevent them (3 marks).
(L3.1: conduct practical investigations to determine the concentrations of solutions and investigate the different ways in which concentrations are measured)
Question 14 ⭐
Calculate the concentration (in mol/L) of the following (6 marks):
Question 15 ⭐⭐⭐
What is the molar concentration of chloride ions in a solution prepared by mixing 100.0 mL of 2.0 M KCl with 50.0 mL of a 1.50 M CaCl2 solution? (5 marks)
(L3.2: manipulate variables and solve problems to calculate concentration, mass or volume using ? = ?/? (molarity formula), dilutions (number of moles before dilution = number of moles of sample after dilution))
Question 16 ⭐⭐⭐
Describe how you would make a standard solution of 0.1M hydrochloric acid and dilute it to 0.02M. Include relevant calculations and safety precautions. (7 marks)
(L3.3: conduct an investigation to make a standard solution and perform a dilution)
Gas Laws
Question 17 ⭐
A student hypothesized that N2 gas in the balloon would deflate upon increased temperature. Will his hypothesis be proven correct or wrong? Explain your answer with reference to Gay-Lussac’s Law and the Ideal Gas Law (3 marks).
(L4.1: conduct investigations and solve problems to determine the relationship between the Ideal Gas Law and Gay-Lussac’s Law (temperature)
Question 18 ⭐⭐⭐
With reference to Boyle’s law, explain the effect of increasing pressure on the volume of a known gas with a constant mass and temperature. How is this represented in the Ideal Gas Law equation? (5 marks)
(L4.1: conduct investigations and solve problems to determine the relationship between the Ideal Gas Law and Boyle’s Law)
Question 19 ⭐⭐
Using the Charles Law, explain the effect of increasing temperature on the molecular characteristics of gases. How is this represented in the Ideal Gas Law equation? (3 marks).
(L4.1: conduct investigations and solve problems to determine the relationship between the Ideal Gas Law and Charles’ Law)
Question 20 ⭐⭐⭐
During an experiment, a student has found that the volume and amount of moles of O2 gas are not proportional to each other. Using the Avogadro’s Law, explain any sources of error and suggest ways to avoid them (4 marks).
(L4.1: conduct investigations and solve problems to determine the relationship between the Ideal Gas Law and Avogadro’s Law)
Worked Solutions
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Question 1
Gravimetric analysis involves the key steps of precipitating, filtering, drying, weighing and calculating. However it is always possible to make mistakes. If the experimental mass is lower than the theoretical mass, an error must have occurred throughout the analysis.
Two examples are; if during filtration, not all of the precipitate is filtered out and some remains in solution, and if scales are not tared correctly or not all of the sample is transferred to the weight plate.
Question 2
Gases have more energy compared to liquids and solids. Therefore if one of the products is a gas, the particles will be moving around more causing an expansion of the volume. If this is hard to visualise, imagine placing the reaction inside a sealed syringe. As the reaction proceeds and gas is formed, the stopper will be pushed out more by the formation of the gas, causing the volume to increase.
Question 3
To investigate the law of conservation of mass focusing on the volume of gases, it is essential that the reactants are carefully measured out and the system is set up to prevent gas escaping. A sealed reaction vessel such as a conical flask with a stopper and attaching a gas syringe to the flask can be used.
Let’s use the reaction between hydrochloric acid and sodium bicarbonate for this investigation. By weighing all the reactants, and the apparatus, we have the required masses for calculations further on. Adding a known volume of acid whilst ensuring the system is sealed allows the reaction to commence without losing any initial gas. Once the reaction has reached completion, the gas formed (carbon dioxide in this case) can be collected and measured in the syringe.
Looking at the system it may seem like the mass has decreased as the solid is “gone”. However, by weighing the sealed system we can show that no mass has been lost. This is a great example demonstrating the law of conservation of mass.
Question 4
💡 Tip: Create a count system using dashes for each element under both the reactants and products side to help you keep track of where you’re at.
Question 5
Question 6
a) To experimentally determine the molar mass of magnesium, a reaction with solid magnesium and hydrochloric acid can be used.
First, accurately weigh the mass of the magnesium ribbon (around 10-15 cm), ensuring that it is clean and any oxidised magnesium has been removed.
Next, pour 50 mL of 1.0M HCl into the conical flask, attaching it to the gas collection apparatus (such as a gas syringe).
Add the magnesium strip, and immediately seal the flask, starting a stopwatch.
The reaction must proceed to completion, ensuring all magnesium has been reacted, and record the total volume of the hydrogen gas that has been produced and collected.
b)
Question 7
The ideal gas law allows us to find the moles of CO2 gas found in the volume occupied using SLC. From there, the moles can be used with the known mass to find an approximate molar mass. As seen in the calculations above, the experimental molar mass is close to the theoretical (44.01 g/mol).
Question 8
Record the mass of Na2CO3and add an excess volume of HCl (recording the volume). Combine the two in a sealed conical flask, connected to a gas syringe. Once the reaction has reached completion the volume of CO2 gas can be recorded.
Using the equation, there is a 1:1 reaction between Na2CO3and CO2, this is what we’re wanting to experimentally confirm. Calculating the moles of Na2CO3using the recorded mass, the expected volume of CO2 can be calculated using SLC conditions. Compare the expected volume to the theoretical volume to show that the reaction proceeds in whole-number ratios.
Question 9
Question 10
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Question 11
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Question 12
Gravimetric analysis relies on the formation of a precipitate. By reacting NaOH with excess CuSO4 a paleblue precipitate will form. Next, the mixture must be filtered to separate the precipitate from the solution, where the precipitate will be washed with distilled water to remove excess reactants.
By drying and weighing the precipitate all water is removed and an accurate mass can be recorded. Stoichiometric calculations can be used to determine the moles of the precipitate, and relating this back to the unknown solution of NaOH, the concentration can be calculated.
Question 13
Three reasons as to why the calculated concentration may be less than the theoretical concentration can occur during the following aspects of gravimetric analysis.
Incomplete precipitation results in not all of the analyte forming a precipitate reducing in a lower mass of precipitate.
During filtration or washing some of the precipitate may be lost or not properly isolated, resulting in a lower mass when weighing.
Lastly, if the precipitate is contaminated and the contaminations are then removed during drying, the precipitate may lose mass.
All these possible reasons result in a lower mass of the precipitate, which will result in a concentration lower than expected.
Question 14
Question 15
💡 Tip: First use the dissociation reactions for each solution separately, before finding the moles of Cl– ions present in each solution.
Question 16
💡 Tip: To create a standard solution you need to start with the concentrated solution. In this case, assume that concentrated HCl is 12 mol/L.
Concentrated hydrochloric acid is often a 12M solution. This is the starting point for creating a 0.1M standard HCl solution. Using C1V1 = C2V2 the amount of 12M HCl required can be calculated as shown below.
This calculation shows that 2.5 mL of 12M HCl, diluted to 300 mL will create the desired 0.1M HCl standard solution.
To dilute this to a 0.02M solution of HCl, the same equation can be used again as shown below.
The final solutionof 0.02M can be achieved by diluting the 0.1M standard to 1.5 L.
As HCl is a strong acid, safety precautions must be followed, especially when handling the concentrated solution. Wearing safety goggles and gloves, as well as ensuring skin is covered (enclosed shoes, long sleeves etc) promote safety throughout this procedure.
Question 17
Question 18
💡 Tip: What relationship does each gas law focus on, and how could this be observed in a gas syringe?
Boyle’s Law states that volume will increase when pressure decreases and vice versa (at a constant temperature). When the gas in the gas syringe is given more “space” (volume) to occupy, the pressure will decrease, however when the volume is decreased by pushing the plunger further in, the pressure will increase as the gaseous particles have less space.
Gay-Lussac’s Law states that temperature increase will increase the pressure if the volume is kept constant. By heating the gas syringe we are increasing the movement of the gaseous particles. If the plunger is fixed, maintaining a constant volume, the pressure will increase.
However, if the plunger is not fixed, the increase in temperature will cause the plunger to be “pushed” further out creating more volume for the gas to occupy due to the increase in pressure. This relates to Charles’ Law which states that hotter gases expand freely if the pressure is kept constant.
(Here is a visual just for reference)
Question 19
At a molecular level when temperature increases, the particles have a greater kinetic energy resulting in faster motion leading to an expansion of volume (if the pressure is held constant).
The relationship established by Charles’ Law that V ∝T, is reflected in the ideal gas law PV=nRT where when pressure (P) and moles (n) are constant, the only variables affecting the answer are volume (V) and temperature (T).
Question 20
Avogadro’s Law states that under the same conditions of temperature and pressure, equal volumes of gases have the same number of moles. If the volume of O2 gas is not proportional to the moles then this is due to deviations in temperature and/or pressure. This means that if the temperature and pressure are not held constant throughout the experiment, it will result in disproportionate values for volume and moles.
By using a controlled environment with efficient temperature regulation and constant pressure, the student can ensure that the proportionality of volume and moles is upheld.
And that wraps up our Practice Questions for Year 11 Chemistry Module 2: Introduction to Quantitative Chemistry – good luck!
You can have a go at our HSC Chemistry practice questions for other modules below:
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Miriam Nelson graduated in 2021 with a Band 6 in Chemistry and is currently studying a Bachelor of Science with a Masters in Secondary Education. However, Miriam has never really left high school, having supported students in Chemistry and Maths for 3+ years while also teaching at a school in Sydney’s CBD. If she’s not studying or learning something new you can find her hunched up over her latest project or at the beach.
Kate Lynn Law (Co-author) graduated in 2017 with an all rounders HSC award and an ATAR of 97.65. Passionate about mentoring, she enjoys working with high school students to improve their academic, work and life skills in preparation for the HSC and what comes next. An avid blogger, Kate had administrated a creative writing page for over 2000 people since 2013, writing to an international audience since her early teenage years.