HSC Chemistry Module 6: Acid-Base Reactions Practice Questions

Question 6 ⭐⭐⭐

Justify the continued use of the Arrhenius definition of acids and bases, despite the development of the more sophisticated Brønsted–Lowry definition. (4 marks)

(L1.6: explore the changes in definitions and models of an acid and a base over time to explain the limitations of each model, including but not limited to: – Arrhenius’ theory – Brønsted-Lowry theory)

Using Brønsted-Lowry Theory

Question 7 ⭐

Complete the table below (3 marks) 

Indicator	Colour in Acid	Colour in Base
Methyl Orange
Bromothymol blue
Phenophathalein

(L2.1: conduct a practical investigation to measure the pH of a range of acids and bases)

Question 8 ⭐⭐

A solution was made by mixing 75.00 mL of 0.120 mol L-¹ hydrochloric acid with 25.00 mL of 0.200 mol L-¹ sodium hydroxide. (4 marks)

What is the pH of the solution?

(L2.2: calculate pH, pOH, hydrogen ion concentration ([H+ ]) and hydroxide ion concentration ([OH– ]) for a range of solutions) 

Question 9 ⭐

Using your understanding of weak and strong acids, explain how the pH of 0.1M of hydrochloric acid would differ from 0.1M of phosphoric acid. (3 marks) 

(L2.3: conduct an investigation to demonstrate the use of pH to indicate the differences between the strength of acids and bases)

Question 10 ⭐⭐

Using ionic equations, explain the amphiprotic nature of sodium hydrogen carbonate or potassium dihydrogen phosphate according to the Bronsted-Lowry definition. (4 marks) 

(L2.4: write ionic equations to represent the dissociation of acids and bases in water, conjugate acid/base pairs in solution and amphiprotic nature of some salts, for example: – sodium hydrogen carbonate – potassium dihydrogen phosphate)

Question 11 ⭐⭐⭐

Referencing the model you have used in class to communicate differences between strong and weak acids, explain why strong acids have a lower pH compared to the weak acids of the same concentration. (4 marks) 

(L2.5: construct models and/or animations to communicate the differences between strong, weak, concentrated and dilute acids and bases)

Question 12 ⭐⭐⭐🔥

500mL of 0.1M hydrochloric acid has been diluted to a volume of 1.25L. The resultant diluted solution was then mixed with 650mL of 0.1M NaOH. Find the pH of the solution. (3 marks) 

(L2.6: calculate the pH of the resultant solution when solutions of acids and/or bases are diluted or mixed)

Quantitative Chemistry

Question 13 ⭐⭐⭐ 🔥

The flowchart shown outlines the sequence of steps used to determine the concentration of an unknown hydrochloric acid solution.

Describe steps A, B and C including correct techniques, equipment and appropriate calculations. Determine the concentration of the hydrochloric acid. (8 marks)

(L3.1: conduct practical investigations to analyse the concentration of an unknown acid or base by titration) 

Question 14 ⭐⭐

The graph shows changes in pH for the titrations of equal volumes of solutions of two monoprotic acids, Acid 1  and Acid 2.

a) Explain the difference between Acid 1 and Acid 2 in terms of their relative strengths and concentrations. (3 marks)

b) Why would phenolphthalein be a suitable indicator for both titrations? (1 mark)

(L3.2: investigate titration curves and conductivity graphs to analyse data to indicate characteristic reaction profiles, for example: 

– strong acid/strong base 

– strong acid/weak base 

– weak acid/strong base)

Question 15 ⭐

In class, you have modelled the neutralization of strong and weak acids and bases using different medias. Explain how the use of scientific models can enhance our understanding of reactions on a molecular scale? (4 marks)

(L3.3: model neutralisation of strong and weak acids and bases using a variety of media)

Question 16 ⭐

A monoprotic acid solution with a concentration of 0.01M has a pH of 4.8 at 25 degrees. Calculate the value of Ka and state whether this acid is likely to be a strong acid or weak acid? (3 marks) 

(L3.4: calculate and apply the dissociation constant (Ka) and pKa (pKa = -log10 (Ka)) to determine the difference between strong and weak acids)

Question 17 ⭐⭐

Outline 3 ways acids or bases are used within the industry and Aboriginal and Torres Strait Islander people. (3 marks) 

(L3.5: explore acid/base analysis techniques that are applied: – in industries – by Aboriginal and Torres Strait Islander Peoples – using digital probes and instruments)

Question 18 ⭐

Sam wants to investigate whether the following items in his household are acidic or basic: Juice, detergent and water. Outline a method he can carry out to test its acidity/ basicity. (3 marks) 

(L3.6: conduct a chemical analysis of a common household substance for its acidity or basicity, for example: – soft drink – wine – juice – medicine)

Question 19 ⭐⭐⭐🔥

Using your understanding of equilibrium, explain how a named buffer works when an acid or a base is added. (4 marks) 

(L3.7: conduct a practical investigation to prepare a buffer and demonstrate its properties)

Question 20 ⭐⭐⭐

Explain the role of the conjugate acid/base pair, H₂PO₄- /HPO₄²-, in maintaining the pH of living cells. Include chemical equations in your answer.

(L3.8: describe the importance of buffers in natural systems) 

Whilst dilute sulfuric acid may be seen as safer as it is less hazardous, it reacts readily with metals. This poses a great risk of corrosion and leaks due to its higher levels of ionisation. Meanwhile, concentrated sulfuric acid can be transported in smaller quantities and does not corrode metals. The equation below indicates their reactivity with water. Dilute sulfuric acid will react fully, whereas concentrated sulfuric acid has fewer free H₃O^– ions due to less availability in water. Therefore, whilst dilute sulfuric acid is better for acid/base reactions, transportation of concentrated sulfuric acid is safer and more efficient.

Question 2

Red cabbage indicator can be used to identify whether a solution is acidic, neutral or basic. However it is only a valid choice of indicator if used correctly. Colours can be misleading if the cabbage extract is too diluted or oxidized, and the gradual colour change makes determining the pH challenging. Furthermore it does not have as precise of a pH indicator compared to the universal indicator or a pH meter. Red cabbage indicator can provide a rough pH estimate and works best for strong acids or bases as the colour differences are more obvious. Therefore, red cabbage can be a valid and accurate indicator if it is used with strong substances to get an estimated pH range.

Question 3

Question 4

💡TIP: Think about the pharmaceutical industry and the food/agriculture industries

Neutralization reactions are significant in ensuring safety, maintaining pH and even to produce the products from these reactions. Indigestion is a common ailment experienced by many people and antacids are a tablet that you can take to neutralise the excess acid produced by the stomach, reducing heartburn or acid reflux. Similarly, bee and wasp stings can be quite painful or annoying, but this is treated by neutralising the reaction, Whilst these two examples may not seem crucial, millions of people all around the world benefit from the available remedies. In industry, areas like swimming pools and soil treatment in agriculture requires the pH to be carefully regulated and maintained. The best way to maintain the pH is to have substances on hand to assist in bringing the pH back to the desired level. In agriculture, acidic soil affects plant growth which would be detrimental to food all over the world. If a swimming pool is too acidic it can irritate the skin and affect other areas negatively. Hence neutralisation reactions are crucial in industry and exceptionally beneficial in everyday life.

Question 5

The experimental value of enthalpy of neutralisation is unlikely to match the theoretical value. This is largely due to the difference in equipment available in schools compared to industrial labs. The equipment in schools is unlikely to be perfectly insulated, causing heat to escape and hence affecting the overall change in temperature. Further possibilities for the difference in enthalpy of neutralisation values is human error throughout the experiment or faulty equipment.

Question 6

💡TIP: Start with definitions for each and how they differ.

The Bronsted-Lowry definition of acids and bases, states that acids are proton donors and bases are proton acceptors. This definition is more advanced as it encompasses non-aqueous reactions and more complex substances like ammonia (NH₃). The Arrhenius definition states acids are substances producing H⁺ ions and bases as producing OH^– ions. The Arrhenius definition is less accurate, but easier to understand and can be applied to many common acids and bases.

Question 7

Question 8

Question 9

Strong acids such as hydrochloric acids readily donate protons (H⁺), resulting in a much higher concentration of H⁺ ions. Meanwhile, weak acids, such as phosphoric acid, are more reluctant to donate their protons, yielding a lower concentration of H⁺ ions. As pH calculations rely on the concentration of H⁺ ions present, hydrochloric and phosphoric acids will have different pH values due to the difference in abundance of H⁺ ions, despite both being acids.

Question 10

Amphiprotic substances are substances that can react as both acids or bases depending on the reaction conditions. For example, sodium hydrogen carbonate (or sodium bicarbonate) is an amphiprotic substance. It is able to both accept or donate a proton (H⁺) as shown below.

Question 11

Whilst pH is determined by the concentration of hydronium ions (or protons) present, the ability for substances to donate these ions will determine the concentration found in solution. The biggest difference between strong and weak acids is that strong acids willingly donate H₃O⁺ ions, resulting in a higher concentration found in solution. Meanwhile, weak acids are less willing to donate the H₃O⁺ ions, leading to a lower concentration of them being present. The concentration of these ions in solution determines the pH and the higher the concentration, the lower the pH. Therefore, even if the concentration of strong and weak acids are the same, the concentration of H₃O⁺ ions found in solution will be different.

Question 12

💡TIP: Don’t get the pOH and the pH mixed up!

Question 13

💡TIP:Visualise yourself doing this practical, what steps did you take?

Step A of this process is creating the standard solution to be used throughout this process. The equipment required is a sensitive balance, beaker and volumetric flask to ensure the solution actually has the desired concentration. Step B is the titration in which an unknown amount of hydrochloric acid is added to a known amount of sodium carbonate until the end point is reached. The end point can be identified using an adequate indicator such as methyl orange, which will be red in acidic solutions and yellow in more basic solution. The laboratory equipment required for a titration includes a burette, pipette, conical flask and stirrer (other equipment will be required to ensure set up of these items). Lastly, step C involves calculations to determine the concentration of the unknown hydrochloric acid solution, as shown below.

Question 14

a) Comparing Acid 1 and 2, Acid 1 starts at a very low pH indicating it is a strong acid. Meanwhile, Acid 2 has a higher pH, implying a weaker acid. However, Acid 1 requires less base for neutralisation compared to Acid 2, suggesting that Acid 1 is a dilute strong acid, whilst Acid 2 is a concentrated weak acid.

b) Phenolphthalein would be a suitable indicator for both as the end point has a pH of around 10, indicating a strong base.

Question 15

Scientific models can enhance understanding of reactions on a molecular scale. With regards to neutralisation reactions, scientific models highlight how the process can be predictable and measurable. The models support understanding regarding the role of ions and ionization, demonstrating pH changes, understanding a titration curve and being able to predict the behaviour of reactions based on the strength of the acid and bases used.

Question 16

Question 17

Aboriginal and Torres Strait Islander people are able to neutralize toxins in plants and treat wounds using their knowledge on acidic and basic substances. The soap tree was used to form a lather to clean or disinfect wounds. Additionally, by chewing the leaves, the basic properties can be used to relieve upset stomachs. Sour fruits (acidic) were used to prevent diseases, thus boosting immune health. The biggest industry using acid-base analysis is the wine industry, where the pH is carefully monitored using pH probes measuring the conductivity.

Question 18

An indicator can be used to test the acidity or alkalinity of the household substances, but this is not as suitable if the substances are coloured. Hence, an indicator would be acceptable to test the pH of water. As the other two items are coloured, a pH probe would be more accurate in determining whether these substances are acidic or basic.

Question 19

💡TIP: Remember that buffers consist of weak acids and their conjugate base (or vice versa)

Buffers consist of an acid and its conjugate base (or a base and its conjugate acid). For example, acetic acid (CH₃COOH) and its conjugate base acetate (CH3COO^–). Weak acids and bases only partially dissociate, which is the key factor to an effective buffer. When a base is added to the buffer solution, the acetic acid donates a proton, forming the conjugate base acetate and thus neutralising the base. When an acid is added to the buffer solution, acetate reacts by accepting the proton and forming acetic acid. Both scenarios rely on Le Chatelier’s principle by re-balancing equilibrium following the change to the system (adding an acid or base). This allows equilibrium to be maintained through the use of the acetic acid/acetate buffer.

Question 20

The pH of blood is required to be in the range of 7.35-7.45, and anything above or below becomes toxic to the body. Such a sensitive range requires the use of an effective buffer such as H₂PO₄^– and HPO₄^2-.

H₂PO₄^– neutralises bases by donating a proton as shown below.
Meanwhile, HPO₄^2- neutralises acids by accepting protons. This dynamic equilibrium allows for the sensitive range of the pH of blood to be maintained.