Textbook 6C.19

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Racquel Fox 2I
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Textbook 6C.19

Postby Racquel Fox 2I » Sun Dec 06, 2020 6:20 pm

I'm a little confused on the answer for part a) and part c). Part a) is comparing HF and HCl and part c) is comparing HBrO2 and HClO2. Why would HCl be stronger than HF, and HClO2 stronger than HBrO2? What is the periodic trend going on?

Akash J 1J
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Re: Textbook 6C.19

Postby Akash J 1J » Sun Dec 06, 2020 6:25 pm

Racquel Fox 3L wrote:I'm a little confused on the answer for part a) and part c). Part a) is comparing HF and HCl and part c) is comparing HBrO2 and HClO2. Why would HCl be stronger than HF, and HClO2 stronger than HBrO2? What is the periodic trend going on?


Remember in the last lecture where Dr. Lavelle said that the weaker the bond between the H-__ , the stronger the acid. The bond between HCl is weaker than HF because the size of the Cl atom is larger than the size of the F atom.

As for HClO2 vs. HBrO2, the Cl atom has a stronger "pulling" ability than Br does (remember, Cl and Br are the central atoms with the H bonded to one of the oxygen, not Cl or Br). This pulling causes O to be more stable I believe, resulting in the anion being more stable, causing HClO2 to disassociate at a higher rate than HBrO2.

IanWheeler3F
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Re: Textbook 6C.19

Postby IanWheeler3F » Sun Dec 06, 2020 6:52 pm

I like to think about it in terms of electronegativity, fluorine is THE most electronegative atom. It is extremely hard to contain it in its diatomic form F2 without it reacting. You will notice that some of the only compounds of noble gases are with fluorine. That is all to say is that when fluorine is attracted to an H+, it does not unbond with it very easily. Since HF will not give up H+ easily it exhibits weak base behavior. A strong acid gives up H+ very easily and since Cl- is so large and H+ is so small the H+ will be given up very easily.

As for the difference between HClO2 and HBrO2, you need to look at the electronegativity and direction of the net dipole moment. Because chlorine is more electronegative there is a smaller difference in EN between Cl-O (0.5) than there is between Br-O (0.7). This means the oxygen has a more negative charge in HBrO2 and by Coulomb's law a more negative charge will exhibit greater attraction so the H+ will be more "stuck" to BrO2- and be less acidic.

Racquel Fox 2I
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Re: Textbook 6C.19

Postby Racquel Fox 2I » Sun Dec 06, 2020 6:54 pm

Akash J 3B wrote:
Racquel Fox 3L wrote:I'm a little confused on the answer for part a) and part c). Part a) is comparing HF and HCl and part c) is comparing HBrO2 and HClO2. Why would HCl be stronger than HF, and HClO2 stronger than HBrO2? What is the periodic trend going on?


Remember in the last lecture where Dr. Lavelle said that the weaker the bond between the H-__ , the stronger the acid. The bond between HCl is weaker than HF because the size of the Cl atom is larger than the size of the F atom.

As for HClO2 vs. HBrO2, the Cl atom has a stronger "pulling" ability than Br does (remember, Cl and Br are the central atoms with the H bonded to one of the oxygen, not Cl or Br). This pulling causes O to be more stable I believe, resulting in the anion being more stable, causing HClO2 to disassociate at a higher rate than HBrO2.


I'm sorry I'm still a little confused as to why for the latter scenario, higher electronegativity creates a stronger acid, but for the first problem, weaker bonds are stronger acids.

Marco Morales 2G
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Re: Textbook 6C.19

Postby Marco Morales 2G » Fri Dec 11, 2020 1:03 pm

HCl is stronger because it has a longer bond than HF, due to trends in the periodic table. Lavelle also mentioned that the longer the A-H bond, the stronger the acid because a property of a strong acid is it gives off protons easily. So if a bond is long, the proton isn't easily held to the nucleus of the central atom, so it'll be given off easily. To make things simpler, just remember that a longer bond length = stronger acid, but i believe it only applies to A-H bonds.

Earl Garrovillo 2L
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Re: Textbook 6C.19

Postby Earl Garrovillo 2L » Fri Dec 11, 2020 4:59 pm

Relative strength of an acid is related to it's ability to lose an H+ ion. So acids where it's very easy to lose a proton are relatively stronger than acids where it is harder to lose an proton. In other words, weaker bonds=stronger acid.

With HF and HCl, the bond between HF is very short because of fluorine's size and high electronegativity. Meanwhile, in HCl's bond is larger because Cl is a significantly larger atom than F and less electronegative so it's easier for HCl to become dissociated.

The HClO3 and HBrO2, it actually seems like the opposite because Cl is more electronegative than HBrO2 but since the H isn't bonded to Cl or Br, the reasoning we had with HCl and HF doesn't apply here and we have to look at the overall dipoles. The more electronegative Cl makes the OH bond more polar and more prone to losing the H+ since polar molecules are more likely to dissociate in water and as a result are stronger acids.

Rose_Malki_3G
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Re: Textbook 6C.19

Postby Rose_Malki_3G » Fri Dec 11, 2020 5:19 pm

HCl is a stronger acid because a strong acid is one that easily loses a proton/hydrogen. Cl has less electronegativity than F so the bond between H and Cl is weaker than that between H and F, making HCl the stronger acid. And for c, HClO2 is stronger because Cl has higher electronegativity than Br so is able to better distribute the charge and make the anion more stable.

Nancy Yao
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Re: Textbook 6C.19

Postby Nancy Yao » Wed Dec 16, 2020 5:30 am

H-Cl bond is longer than H-F bond because Cl has larger atomic radius than F, so it is easier to break. Strong acid loses H+ more easily, and also the resulting anion is stabler. For HClO, the Cl in the Cl-O anion delocalize electron to a larger degree than Br in Br-O, and therefore HClO has greater acidity than HBrO.


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