pH question... HELP PLEASE (1 Viewer)

[punk0rz]

could someone please help me to understand the process of determining the pH of a solution?
is it required that we determine the pH of solutions given their formula's and concentrations?
how does the dissociation of ions of an acid/base influence the process, shouldnt a strong acid such as HCl dissociate completely and so the concentration of the solution can be taken as the concentration of [H+] (as the ammount of molecules which do not dissociate is negligible) while a weak acid such as ethanoic acid which does not dissociate completely will required a different value for concentration of [H+] to the concentration of the solution, as some of the molecules will remain in complete form??
am i rambling complete shit? do i not even need to understand this?

heeellppp im going insane

and thankyou in advance

 

[Frigid]

3.2.1 Define acids as proton donors and describe the ionisation of acids in water.

Acids can be defined as proton donors because they contribute a proton (a hydrogen ion) to water when they ionise. For example:

CH3COOH(aq) + H2O(l) --> H3O+(aq) + CH3COO(aq)
Acetic acid + Water --> Hydronium ion + Acetate ion

3.2.2 Identify acids including acetic (ethanoic), citric (2hydroxypropane1, 2, 3tricarboxylic), hydrochloric and sulfuric acid.

Acetic Acid: (Ethanoic acid) CH3COOH(aq)

Citric Acid: (2hydroxypropane1, 2, 3tricarboxylic acid)
(COOH)CH2CH(OH)(COOH)CH2COOH(aq)

Hydrochloric Acid (Hydrochloric acid) HCl(aq)

Sulfuric Acid (Sulfuric acid)
H2SO4(aq)

3.2.3 Describe the use of the pH scale in comparing acids and bases.
3.2.5 Identify pH as log10[H+] and explain that a change in pH of 1 means a ten-fold change in [H+].

The pH scale is used to compare acids and bases, based on the concentration of [H+] in solution. pH is measured from 0 to 14 (acidic to basic). The pH scale is a logarithmic relationship, the formula being:

pH = log10[H+]

The pH of a base is derived by the following:

pOH = log10[OH]
pH = 14 pOH

where [OH] is the concentration of hydroxide ions in the solution.

3.2.4 Describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and dilute.
3.2.6 Compare the relative strengths of equal concentration of citric, acetic and hydrochloric acid and explain in terms of the degree of ionisation of their molecules.

Acids ionise in water and become proton donors, forming [H+] ions in water. The greater the concentration of [H+], the greater is the strength of the acid. The negative logarithmic relationship only works for strong acids, those which ionise completely in water. A weak acid is an acid which only partially ionises in solution. A concentrated acid can be either strong or weak. Concentration and weakness refer to the level of dilution of the solution using a solute, usually water.
Between hydrochloric, acetic and citric, the strong acid, hydrochloric, is strongest. The weakest acid is the weak acid, acetic, and citric acid is between the strengths of the two.

3.2.7 Describe the difference between a strong and weak acid in terms of an equilibrium between the intact molecule and its ions.

In a strong acid, such as hydrochloric acid, an equilibrium is formed during ionisation:

HCl(aq) + H2O(l) --> H3O+(aq) + Cl(aq)
Hydrochloric acid + Water --> Hydronium ion + Chloride ion

In the equilibrium of the strong acid, the equation completely lies on the right side (near 100% ionisation). The molecule of the strong acid completely ionises.

In a weak acid, such as acetic acid, an equilibrium is formed during ionisation:

CH3COOH(aq) + H2O(l) --> H3O+(aq) + CH3COO(aq)
Acetic acid + Water --> Hydronium ion + Acetate ion

In the equilibrium of the weak acid, the equation lies mostly on the left (partial ionisation). The molecule of the weak acid is in solution with few of its ions.

 

[mon_mon]

jesus, i was going to reply, but i see there's no need. Pretty much sums it up. But just in case....
pH is equal to the -log of your H+ concentration thus you enter '-' then 'log' then the H+ number into your calculator.

 

[toknblackguy]

if they said for instance that the weak acid ethanoic acid dissociates only 15%, and this was reacted again strong NaOH,
you write up the equation - CH3COOH + NaOH <-->CH3COONa + H20, then i the stroichemetric ratio of hydrogen to hydroxy is 1:1. you work out the number of moles assuming complete dissociation, and then muptpliy by 0.15. then you see which one is in excess (in this case OH-) and then the pH of the solution would be 14 + log[OH-]

frigid or someone else
correct me if i'm wrong
cheers

 

[spice girl]

you're wrong...but the actual working out of the pH of ethanoate is too complex for HSC level...it requires things like acid dissociation constants.

in equilibrium chemistry we'd rarely say that smth has 15% dissociation, because the percentage dissociation depends on the concentration of the weak acid / base.

 

[Frigid]

WRONG!

(off the top of head again): in neutralisation reactions, even if one part is weak acid/base, the reaction will still complete itself... why?

think of it this way:

CH3COOH(aq) <==> H+(aq) + CH3COO-(aq)

the equilibrium lies well to the left.

now think, under Mr. LeChats, what would happen if i take one H+ away and combine it with an OH- from NaOH (NaOH is a strong base, assume complete dissociation). under Mr. LeChats, the concentration of H+ will go down, therefore system will drive itself forward to create more H+. this process repeats itself until ALL the OH- available pulls and combines with all the H+ available (in ratio 1:1). then you calculate pH based on whatever's excess, or if the salt formed is basic/acidic (another story, another day...)

therefore you are wrong.

edit: of course, i also might be completely wrong. if i am, i apologise to all you IChO winners out there...

 

[toknblackguy]

hahahaha
i got owned

that certainly explains a lot...note to self...never trust a known bs-artist (ppl from my school may know who i'm talking about)
thanks

 

[mon_mon]

All you need to know is that weak acids will have a higher pH (thus less acidic) than strong acids, which will always have a pH value of 1 for a 0.1 molar soln. providing they are monoprotic. di and triprotic is a whole other barrel of fish that we don't need to know about.

 

[inasero]

ohhh you're treading on thin ice refuting the statements of an ICho winner, but I'm inclined to agree with you since I don't undertsand this higher level Chem (therefore not relevant to syllabus)...

 

[mon_mon]

Umm...i think thats a complement. actually, screw you! Nah, im kidding.

 

[Exeter]

i read somewhere that ph of acetic(ethanoic) acid is 2.9 for a 0.1M soln...and that it is about 99% CH3COOH ions, and 1% the dissociated

 

[mon_mon]

Yup. probably even less for many acids. take citric for example. or even Asprin! it would be about 0.00000000000000001% disociated.