April 25, 2024, 06:29:11 AM
Forum Rules: Read This Before Posting


« previous next »
Pages: [1]   Go Down

Topic: pH of nitric acid  (Read 15080 times)

0 Members and 1 Guest are viewing this topic.

Offline Wil"

  • Regular Member
  • ***
  • Posts: 40
  • Mole Snacks: +0/-3
pH of nitric acid
« on: March 18, 2009, 06:34:54 AM »
I want to ask what's the pH of a 1x10-8M nitric acid(aq)?

I know how to calculate the pH,
pH = -log[HNO3]
which gives you pH=8 in this case.

But nitric acid is an acid, isn't it?? How its greater than 7??

I checked wiki before and I got "nitric acid" doesn't donate H+

So what's so dfferent of HNO3??

Can  anyone tell me more?? I just couldn't find it....

Thanks!
Logged

Offline sjb

  • Global Moderator
  • Sr. Member
  • ***
  • Posts: 3652
  • Mole Snacks: +222/-42
  • Gender: Male
Re: pH of nitric acid
« Reply #1 on: March 18, 2009, 07:35:48 AM »
Quote from: Wil" on March 18, 2009, 06:34:54 AM
I know how to calculate the pH,
pH = -log[HNO3]

No, the definition of pH is -log[H+].

Try again, with this revised information.
Logged

Offline Wil"

  • Regular Member
  • ***
  • Posts: 40
  • Mole Snacks: +0/-3
Re: pH of nitric acid
« Reply #2 on: March 18, 2009, 10:21:02 AM »
but the value won't make any difference, because 1x10-8M HNO3 gives 1x10-8M H3O+

still I got pH=8....
Logged

Offline sjb

  • Global Moderator
  • Sr. Member
  • ***
  • Posts: 3652
  • Mole Snacks: +222/-42
  • Gender: Male
Re: pH of nitric acid
« Reply #3 on: March 18, 2009, 10:27:10 AM »
Why are you concerned that the pH is greater than 7?

What solvent are you measuring the pH in?
Logged

Offline AWK

  • Retired Staff
  • Sr. Member
  • *
  • Posts: 7979
  • Mole Snacks: +555/-93
  • Gender: Male
Re: pH of nitric acid
« Reply #4 on: March 18, 2009, 10:52:22 AM »
Take into account also hydrogen ions from water, and solve quadratic equation.
Logged
AWK

Offline Wil"

  • Regular Member
  • ***
  • Posts: 40
  • Mole Snacks: +0/-3
Re: pH of nitric acid
« Reply #5 on: March 18, 2009, 11:14:31 AM »
Quote from: AWK on March 18, 2009, 10:52:22 AM
Take into account also hydrogen ions from water, and solve quadratic equation.

but why so sudden I have to take the H+ from water into account??

I got it 6.95....
Logged

Offline Borek

  • Mr. pH
  • Administrator
  • Deity Member
  • *
  • Posts: 27663
  • Mole Snacks: +1801/-410
  • Gender: Male
  • I am known to be occasionally wrong.
    • Chembuddy
Re: pH of nitric acid
« Reply #6 on: March 18, 2009, 11:21:01 AM »
Generally speaking you should take them into account always, however, for more concentrated acids they can be safely ignored - their concentraion is just too low to play any role. That's not th ecase when their concentration becomes comparable with concentration of acid.

6.95 doesn't look correct to me.

http://www.chembuddy.com/?left=pH-calculation&right=pH-strong-acid-base
Logged
ChemBuddy chemical calculators - stoichiometry, pH, concentration, buffer preparation, titrations.info

Offline Loyal

  • Regular Member
  • ***
  • Posts: 84
  • Mole Snacks: +7/-5
  • Gender: Male
Re: pH of nitric acid
« Reply #7 on: March 19, 2009, 02:46:38 AM »
Quote from: Wil" on March 18, 2009, 11:14:31 AM
Quote from: AWK on March 18, 2009, 10:52:22 AM
Take into account also hydrogen ions from water, and solve quadratic equation.

but why so sudden I have to take the H+ from water into account??

I got it 6.95....

When you are dealing with any acid concentration that is around 10-5M or lower the H+ produced from water is enough to make an impact or in the case of 10-7M or lower the water actually produces as much or more H+.   
Logged
Chemistry Student(Senior) at WSU
Pages: [1]   Go Up
« previous next »
Sponsored Links