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Thread: Titration Calculations

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    Titration Calculations

    Can anyone help me with these questions, I'm stuck on them

    1. What action would result in an increase of TWO pH units of the solution?
    (a) Diluting 10mL of 0.01mL of 0.01 M HCl(aq) to 40mL
    (b) Diluting 10mL of 0.01M NaOH(aq) to 40mL
    (c) Diluting 10mL of 0.01M HCl(aq) to 1000mL
    (d) Diluting 10mL of 0.01M NaOH(aq) to 1000mL

    2. A 0.1M HCl solution has a pH of 1.0. What volume of water must be added to 90mL of this solution to obtain a final pH of 2.0?
    (a) 10mL
    (b) 180mL
    (c) 810mL
    (d) 900mL

    3. A number of solutions were tested with a conductivity probe attached to a data logger. Which of the following solutions would record the highest conductivity reading?
    (a) 0.01M HCl
    (b) 0.1M HCl
    (c) 0.01M CH3COOH
    (d) 0.1M CH3COOH

    I think the answer for question 3 is B? Because strong acids ionize completely in solution compared to a weak acid, therefore there is more ions flowing around the solution? Also, I think it's B and not A because it has a higher concentration therefore, there are more HCl particles in the solution, which as a result gives high conductivity readings? Can anyone explain this question for me particularly?


    Thanks in advance to whoever explains to me question 1&2 with calculations for the multiple choice!

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    Ancient Orator leehuan's Avatar
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    Re: Titration Calculations

    (Note: NONE of your questions were on titrations)

    1. The formula pH = -log_10[H3O+] tells us that for every increase by 1 in the pH, the concentration of the hydronium ion is decreased TEN-FOLD. So if we increase pH by 2, the concentration is decrease one hundred-fold.
    Hydrochloric acid - strong acid and monoprotic implies [HCl]=[H3O+] - this I'd hope you can pick up by yourself.
    So using C=n/V by inspection if V=10mL goes to V=1000mL, that's increasing the volume 100-fold, and thus decreasing the concentration 100-fold (moles is constant). So the answer is C.

    2. Same principle, but the pH increase is by 1 this time. So we want to decrease the concentration 10-fold, and we do so by increasing the volume 10-fold (C=n/V). V=90mL, so we have to make V=900mL. 900-90=810. The answer is C.

    3. We will first eliminate C and D as acetic acid is a weak acid and does not fully ionise, whereas hydrochloric does.
    Now, if my memory serves me right, this question has one, massive, trap.

    The more concentrated HCl is, the more MOLECULAR HCl IS PRESENT.
    What happens is, especially if you get to this in industrial chemistry with sulfuric acid, the more concentrated an acid is the less it actually ionises in solution. The most concentrated acids are virtually molecular.

    So you actually just have HCl(g) DISSOLVED in the water. You don't truly have H+ and Cl- anymore.

    Because a substance has to be ionised to conduct electricity, if the concentrated acid has less degree of ionisation, it means that it is LESS electrically conductive than the DILUTE one.

    So if my memory serves me right, the answer is actually A.

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    Re: Titration Calculations

    Quote Originally Posted by leehuan View Post
    (Note: NONE of your questions were on titrations)

    1. The formula pH = -log_10[H3O+] tells us that for every increase by 1 in the pH, the concentration of the hydronium ion is decreased TEN-FOLD. So if we increase pH by 2, the concentration is decrease one hundred-fold.
    Hydrochloric acid - strong acid and monoprotic implies [HCl]=[H3O+] - this I'd hope you can pick up by yourself.
    So using C=n/V by inspection if V=10mL goes to V=1000mL, that's increasing the volume 100-fold, and thus decreasing the concentration 100-fold (moles is constant). So the answer is C.

    2. Same principle, but the pH increase is by 1 this time. So we want to decrease the concentration 10-fold, and we do so by increasing the volume 10-fold (C=n/V). V=90mL, so we have to make V=900mL. 900-90=810. The answer is C.

    3. We will first eliminate C and D as acetic acid is a weak acid and does not fully ionise, whereas hydrochloric does.
    Now, if my memory serves me right, this question has one, massive, trap.

    The more concentrated HCl is, the more MOLECULAR HCl IS PRESENT.
    What happens is, especially if you get to this in industrial chemistry with sulfuric acid, the more concentrated an acid is the less it actually ionises in solution. The most concentrated acids are virtually molecular.

    So you actually just have HCl(g) DISSOLVED in the water. You don't truly have H+ and Cl- anymore.

    Because a substance has to be ionised to conduct electricity, if the concentrated acid has less degree of ionisation, it means that it is LESS electrically conductive than the DILUTE one.

    So if my memory serves me right, the answer is actually A.
    Yo @Leehuan, thanks for replying. I understand completely for question 1 &2. However, for question 1, can you explain why it isn't (D)? And for questions 1&2, is there another method to do these questions involving moles, concentration and dilution formulas?

    For question 3: So you're saying the more concentrated an acid, the less it ionises? I don't understand the bit "most concentrated acids are virtually molecular" and "So you actually just have HCl(g) DISSOLVED in the water. You don't truly have H+ and Cl- anymore".

    I studied Acidic Environment thoroughly but I didn't read anywhere that it says, more concentrated acids ionise less. Or is this something I'll study in Industrial Chem?

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    Ancient Orator leehuan's Avatar
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    Re: Titration Calculations

    Quote Originally Posted by RachelGreen View Post
    Yo @Leehuan, thanks for replying. I understand completely for question 1 &2. However, for question 1, can you explain why it isn't (D)? And for questions 1&2, is there another method to do these questions involving moles, concentration and dilution formulas?

    For question 3: So you're saying the more concentrated an acid, the less it ionises? I don't understand the bit "most concentrated acids are virtually molecular" and "So you actually just have HCl(g) DISSOLVED in the water. You don't truly have H+ and Cl- anymore".

    I studied Acidic Environment thoroughly but I didn't read anywhere that it says, more concentrated acids ionise less. Or is this something I'll study in Industrial Chem?
    Diluting NaOH (a base) would decrease the pH back down.
    ______________

    I learnt about it in industrial chemistry tbh so...

    But I mean, if you think about it it kinda makes sense. When the acid is so concentrated, you don't even have much water anymore. How can you ionise without water being present?
    Last edited by leehuan; 29 Feb 2016 at 8:46 PM.

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    Re: Titration Calculations

    Thanks @leehuan, that's really interesting. I'll keep that in mind when doing exam questions! Thanks for helping, appreciate it

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