AI and chemistry

 How good are AI generated responses to typical chemistry questions?

No better than the source material used for the answer ... which can include common misconceptions.

One such question is "What is the pH of water?". To a chemist this is a nonsense question (comparable to asking what colour an apple is or what shape a building is) yet it is often asked in school science classrooms around the world. 

I asked ChatGPT "What is the pH of water?" This was the answer ...

The pH of pure water at 25 degrees Celsius (77 degrees Fahrenheit) is considered neutral and is defined as 7. A pH value below 7 indicates acidity, while a pH value above 7 indicates alkalinity. It's important to note that the pH of water can be influenced by dissolved substances, and natural water sources may have varying pH levels due to the presence of minerals and other compounds.

I asked perplexity.ai "What is the pH of water?" This was the answer ...

The pH of water is 7, which is considered neutral. A pH of 7 is neither acidic nor basic. It is important to note that the pH of pure water can change due to the presence of impurities or the absorption of carbon dioxide from the air, which can make it slightly acidic. However, under normal conditions, the pH of pure water is 7, indicating neutrality

2

Surprise, surprise, both answers are inaccurate and riddled with common misconceptions (although, to its credit, perplexity.ai does provide sources, but to its discredit it does not refer to a specific temperature which is critical!). ChatGPT emphasizes "pure water", while perplexity.ai is content with just "water". To a chemist, both terms are the same, "water" is a pure substance. If water has substances dissolved in it it is no longer "water" it is an "aqueous solution". Perplexity.ai's referral to "normal conditions" is baffling, perhaps it refers to standard conditions for aqueous solutions? After these issues, both AIs seem to perpetuate the misconception that acidity, basicity (alkalinity) and neutrality are defined by a pH value. This is not true. They are two entirely different concepts (although both are based on an Arrhenius description of aqueous solutions). pH is a measure of aqueous hydrogen ion concentration, independent of whether the aqueous solution has been designated acidic, basic or neutral. Acidic, basic, neutral are terms to describe a the relative amounts of hydrogen ions and hydroxide ions in an aqueous solution.

So, let's start by talking about the neutrality of "pure water". Pure water is always neutral (regardless of its pH). This is because the concentration of acidic protons in the water, H⁺(aq) or H₃O⁺(aq), is always equal to the concentration of basic hydroxide ions in the water, OH^-(aq), that is [H⁺(aq)] = [OH^-(aq)]

An acidic aqueous solution is one in which the concentration of acidic protons is greater than the concentration of basic hydroxide ions, that is, [H⁺(aq)] > [OH^-(aq)]

A basic aqueous solution is one in which the concentration of acidic protons is less than the concentration of basic hydroxide ions, that is, [H⁺(aq)] < [OH^-(aq)] 

So "pure water" is always neutral, but what is its pH? pH is a measure of the concentration of aqueous hydrogen ions in solution, the greater the concentration of hydrogen ions in aqueous solution, the lower the pH. The concentration of these hydrogen ions is very much dependent on the water temperature. The pH of pure water at 10^oC is approximately 7.27 but the water is still described as neutral because [H⁺(aq)] = [OH^-(aq)]. The pH of pure water at 50^oC is about 6.63 but the water is still neutral because [H⁺(aq)] = [OH^-(aq)].

 The pH of water at 25^oC is approximately 7 and is based, NOT on it being defined this way, but on a calculation of the hydrogen ion concentration of water at this temperature.

So, what happens to pH if other substances are dissolved in water to form aqueous solutions at 25^oC? Using an Arrhenius definition of acids and bases (and a pH of 7 which has 1 significant figure but even that digit is uncertain), if an aqueous solution has pH < 7  at 25^oC it could be described as acidic, but if the temperature is greater than 25^oC the solution may be neutral, it may even be basic, we simply do not have enough information to decide. Similarly, If an aqueous solution has pH > 7  at 25^oC it could be described as basic, but if the temperature is less than 25^oC it is not a justifiable conclusion.

It is important to note that if there are dissolved "substances", including "impurities", "minerals" or "carbon dioxide", then the "water" is not "pure water", it is an aqueous solution, and the AI discussions are not relevant. Furthermore, some solutes dissolve in water at 25^oC and increase its pH  (aqueous solution of sodium hydroxide), some decrease its pH  (aqueous solution of hydrogen chloride), and some may have no effect at all on its pH (aqueous solution of sodium chloride).

There is no simple, correct answer for the question "What is the pH of water?". The question does not contain enough information in order for it to be answered. IF the AIs simply stated that the pH of water at 25^oC is approximately 7, that would be a pretty good response, however, by providing additional information in the answer they are repeating, and encouraging, common misconceptions about the nature of "pure water", "pH", "acidity", "basicity" (or "alkalinity") and "neutrality". 

For a discussion on why pure water is always neutral and why its pH varies, see https://www.ausetute.com.au/kw.html

For a discussion on what determines whether an aqueous solution is acidic, basic or neutral, see https://www.ausetute.com.au/abneutral.html

Chemical Equations for Equilibrium Reactions

 

When do you use single arrow → in a chemical equation and when do you use a double-headed arrow ⇌ in a chemical equation?

Go to https://www.ausetute.com.au/equilibriumeqtn.html to find out!

 

AUS-e-TUTE Members should log-in to use the new Writing Chemical Equations for Equilibrium Reactions tutorial, game, test and exam with worked solutions.

Coming to terms with chemical equilibrium

In order to understand chemical reactions, you are going to have to come to terms with an important concept, that is, chemical equilibrium.

In AUS-e-TUTE's new set of resources (tutorial, game, test and exam questions with worked solutions) we look at how the type of system (open or closed) and the type of reaction (reversible or irreversible) determines whether a system can achieve a state of balance otherwise known as chemical equilibrium.

AUS-e-TUTE members should log-in to use these new resources.

If you are not an AUS-e-TUTE members you can view the tutorial for free at 

https://www.ausetute.com.au/equilibriumsystem.html

Chemistry of Tooth Decay and Prevention

Tooth decay is a major health concern.

For more than 50 years we have been adding fluoride to our water supplies and toothpaste to help prevent tooth decay.

It is estimated that this fluoridation of drinking water alone has resulted in a 25% decrease of tooth decay in children and adults.

What causes tooth decay and how does fluoride prevent tooth decay and cavities?

 Learn about the chemistry of tooth decay and how to prevent it in the March 2021 issue of AUS-e-NEWS. Subscribe to AUS-e-NEWS for free at https://www.ausetute.com.au/ausenews.html

Emergency Oxygen

Have you ever sat on a plane listening to the flight attendants give the safety demonstration,

"In case of cabin depressurization, oxygen masks will drop automatically..."

and wondered about where, and how, this emergency oxygen is stored?

 Read this edition of AUS-e-NEWS to find out! 

You can subscribe to AUS-e-NEWS at https://www.ausetute.com.au/ausenews.html

pH of Aqueous Salt Solutions

If you titrate acetic acid with sodium hydroxide, what is the pH at the equivalence point?

If you dissolve ammonium chloride in water, what is the pH of the solution?

To answer these questions you will have to calculate the pH of an aqueous salt solution!

AUS-e-TUTE has just added new resources to help you do just that!

Members should log-in to use the new tutorial, game, test, drill, worksheet wizard, and problem solving template.

If you are not an AUS-e-TUTE Member, you can access the "free-to-view" tutorial at https://www.ausetute.com.au/phsaltcalcs.html

Equilibrium Constants are Affected by Temperature Changes

Question:
How would the value of the equilibrium constant, k, be affected if the reaction was heated by 1000°c?
(a) if the reaction was endothermic
(b) if the reaction was exothermic

Answer:
(a)  For an endothermic reaction, an increase in temperature will increase the value of the equilibrium constant.
(b) For an exothermic reaction, an increase in temperature will decrease the value of the equilibrium constant.

Explanation: 

Go to https://www.ausetute.com.au/ktemperature.html for the tutorial. 

AUS-e-TUTE Members should log-in to use the game, test and exam with worked solutions on this topic.

pK

You've probably heard of pH, maybe even pOH, but what is pK?
Does this sound like brand of chewing gum to you?
Or does that "p" mean something?
Time to find out, with AUS-e-TUTE's new pK resources!

https://www.ausetute.com.au/pk.html

pH of Weak Acid After Dilution

What happens to the pH of a solution of weak acid, such as acetic acid, if you dilute it by adding water to it?
How much water would you have to add to 100 mL of 0.20 mol L^-1 acetic acid in order to make a solution with a pH of 3.0 ?
What is the pH of 1 × 10^-11 mol L^-1 acetic acid?
AUS-e-TUTE can help you answer these questions with our new resources.
AUS-TUTE Members should log-in to access the new "pH of Weak Acid After Dilution" tutorial, game and test (with worked solutions).
A "free-to-view" tutorial is currently available for non-members at:
https://www.ausetute.com.au/phdilwa.html

Gibbs Free Energy and Spontaneity of Reactions

Whether or not a chemical reaction proceeds in a particular direction depends on a balance between the enthalpy of the system and its entropy.
Gibbs Free Energy allows us to quantify this relationship, and determine whether a particular reaction will be spontaneous.
AUS-e-TUTE members can now access a new tutorial, game, test and exam on this topic.
Not a member?
There is a "free-to-view" tutorial currently available at http://www.ausetute.com.au/freeenergy.html

Acid-Base Equilibria and Beginner Teachers

Cathy is a Year 12 student in a Brisbane school (capital city of Queensland, Australia). Last week she had a chemistry exam, the topic was equilibrium which included acids, bases and acid-base titrations. I met her on her way to school on the morning of this exam. When I asked her if she felt confident about her exam, I was horrified by her response,
"Sort of. I've got a new teacher this year and she's not any better than the one I had last year. We did an experiment, she said it took too long so we didn't do any more."
Seasoned teachers are used to:

• student claims that their teacher is "no good" (especially if the student is performing poorly)
• student exaggeration (only 1 experiment in a whole term of equilibrium, surely not!)

"Oh come on", I said in my best 'you're kidding me' voice, "you studied equilibrium for a whole term and only did one experiment?"

"Yeah", she confirmed, "we watched a video on titration though".

"Didn't you do a titration experiment?"

"Nah. She said we didn't have time."

Even allowing for the possibility of student exaggeration, the thought that you would play a video showing someone else performing a titration rather than giving your own students the opportunity to carry out even a simple titration, is, quite frankly, appalling. 

From Cathy's description of the teacher I assume this is the teacher's first year out teaching (beginner teacher). Reflecting on my own first year of teaching (a long, long time ago), I remember struggling to meet the requirements of the chemistry syllabus in the time-frame allowed, and, I also remember that while teaching the techniques of titration was time consuming, the students learned more in a few practical sessions than they learned during the whole of the preceding theoretical lessons, and once they have mastered the techniques they can be put to use in real-world problems.

A typical sequence that is often taught for (monoprotic) acid-base equilibria assumes prior knowledge of solutions, concentration (molarity) and equilibrium concepts and calculations (including self-dissociation of water, K_w ):

1. What are the properties of acids and bases?
2. How do we define an acid and a base?
3. What is meant by the terms "strong acid"  and "strong base"?
4. How do we measure the strength of an acid or a base (pH scale)?
5. What happens when you add an (Arrhenius) acid to a (Arrhenius) base (neutralisation)?
6. How much (Arrhenius) acid do we need to add to a known amount of (Arrhenius) base in order to neutralise it (acidic, basic, neutral solutions)?
7. Discussion of titration techniques, including preparation of a standard solution.
8. Performing a strong acid - strong base titration.
9. Using the results of the experiment to calculate the concentration of the unknown acid or base.
10. Perform calculations for each 1 mL addition of strong acid to strong base in the titration experiment and graph the results (strong acid - strong base titration curve)
11. Discussion of weak acids (K_a).
12. Discussion of other acid-base reactions (proton transfer reactions) and other titration curves
13. What indicator should you use for a particular acid-base reaction? (optional, how does an indicator work)
14. Titration of weak acid - strong base (such as determination of acetic acid in vinegar)

If you see your students 4 or 5 times a week, this teaching program for acid-base equilibria will take about 4 weeks using a traditional, structured approach. If you have the luxury of being able to time your practical work so that it occurs in the correct sequence, and take time to link the practical work to the theoretical concepts, your students have a good chance of understanding and being able apply the concepts to unfamiliar problems.

If you don't do any experimental work, you could probably bowl it over in 2 to 3 of weeks, and be faced with a lot of bored students wondering why they ever took a course in chemistry.

If you take a student-centered constructivist approach (for example, start with the questions  like "what gives vinegar its tangy taste?",  "if acids are corrosive and burn skin, how come you can drink vinegar?",  "how can you measure the strength of an acid?", "how can we determine which brand of vinegar has the greatest concentration of acetic acid?"), be prepared to add another week (unless you give the students a lot of reading/research for homework). The benefits, however, are enormous. Your students are more likely to be engaged with the content and "on task", they will have to be able to justify decisions they make in order to design and perform experiments thereby linking concepts and practical work, and because they "invest" in the whole learning process they are more likely to be apply the understanding and knowledge gained to other problems.

So, if you are new to teaching acid-base equilibria, here a few suggestions:

1. Even if you firmly believe that constuctivist approaches to teaching are the most effective way to teach chemistry, be prepared to spend your first year of teaching chemistry taking a more traditional approach, using guided questioning to lead students towards the experiment(s) you need them to do (syllabus requirements) while still giving them "ownership" of the experiment and its results. Keep a list of the misconceptions you come across when you teach, this will help you be better prepared for next year. As you feel more confident in your ability to meet the syllabus requirements within the time you have, and you have a better feel for the misconceptions you will meet, you can start "loosening your hold" and give more time to truly constructivist approaches.
2. Let the students do as much practical work as possible (students not only need to be exposed the practical techniques of chemistry, they need to do the experiments in order to fully appreciate the significance of what you are trying to teach them). You also need to devote time to discussing the results of their experiments with them, and reinforcing the concepts, calculations, techniques etc involved. 
3. Spend time discussing the self-dissociation of water (that is, it is a lesson in its own right, not just a passing reference before you discuss acid-dissociation). Students will have been exposed to an "acids and bases" topic sometime between Years 7 and 10, but even so, many of them may still think that an acid has a pH less than 7, a base has a pH greater than 7, and that a neutral substance has a pH of 7. Believe me, it can be an uphill struggle to separate the two concepts of "acid, base, neutral" from the concept of "pH" in a student's mind (and if you don't believe me, think about the number of times you have seen/heard advertisements for products which talk about the "neutral pH" of skin/hair etc). If the students do not have a good grasp of the self-dissociation of water then they will not understand the pH of  aqueous solutions. (And a word of caution, just because a student can calculate the pH of an aqueous solution of base at 25^oC, it doesn't mean they understand the relevance of pH + pOH = 14, or [H⁺][OH^-] = 10^-14, and if you want to test this statement, ask you students to calculate the pH of 0.001 M NaOH(aq) at 50^oC, or ask them to find the pH of 0.01 moles of HCl(g) dissolved in 1 L of ethanol and see what happens, because the chances are they will simply do a pH + pOH = 14 calculation without even thinking about it!)
4. Spend time making the distinction between "strong acids", "weak acids", "dilute aqueous solutions of acids" and "concentrated aqueous solutions of acids" (similarly for bases) because once again, you are likely to have an uphill struggle to separate the two concepts "strength of an acid/base" and "concentration of an acid/base". Remember, they have already been exposed to statements such as, "I need a cup of strong coffee", or, "this cordial is a bit strong" which, in chemical terms should be "I need a cup of concentrated aqueous solution of coffee (or cordial as the case may be)". On the other hand, they have also been exposed to ads which say things like "concentrated laundry detergent" which is a slightly more appropriate use of the technical term "concentrated" (although I do remember one example that used "concentrated laundry liquid" which introduces the other problem of the loose usage of the word "liquid" instead of "solution"). One way to do this is to give each pair of students a bottle of acetic acid labelled with its concentration, and have them measure its pH with a pH meter. Also provide them with volumetric flask of HCl(aq) of known concentration (say 0.1 M) and have them measure its pH, then have them perform sequential 1:10 dilutions and measure the pH at each stage say they can see that pH is dependent on the concentration of the strong acid and that you can reach a point at which the pH, and therefore the concentration, of a strong acid is the same, and even greater than, the concentration of an aqueous solution of weak acid. When you tabulate the class results and ask them for an explanation be prepared for many of them to believe you somehow "tricked them", it can take time for them to break the strength/concentration misconception and replace it with a more appropriate separation of the two concepts. If the students do this activity themselves, it will easily take a lesson, if you do it as a demonstration it will take about 10 minutes, BUT, it is better for the students to do it themselves partly because it reduces the instances of "there must be a trick in this" thinking, but mostly because they can see the pH change with the concentration and they are going to have to justify that all the way to the point at which the pH of the strong acid is  greater than the pH of a weak acid.
5. Related to point 2 is the common misconception students may have that when you add an acid to a base you end up with a neutral solution that has a pH of 7. Personally, I think the best way to deal with this is to let the students work it out for themselves before you attempt to explain it to them. For example. give each pair of students a bottle of methyl orange indicator (you will need a fair degree of tolerance in establishing the end-point so don't use a pH meter) and a conical flask and have them add a 10.00 mL aliquot of standardised 0.1 M NaOH to the flask and record what happens to the indicator colour. Have them calculate the moles of hydroxide ions in the flask, as well as calculate the pH of the solution (so they are convinced the indicator is giving a true reading). Then give each pair a 100 mL volumetric flask containing 0.1 M monoprotic acid (some will get a strong acid such as HCl(aq), some will get a weak acid such as acetic acid). . Have them calculate the volume of acid they will need to add to the NaOH(aq) to neutralise it. Have them add this volume (straight from the pipette to the flask), give it a swirl, and record the colour of the indicator. Tabulate the results on the board (yellow vs red). Ask them why some changed colour and some didn't (be prepared to let the students with weak acids try adding more acid, many will believe they made a mistake in the calculations or in adding the solutions), if the students do not come to the realisation that only students with strong acids got a colour change at neutralisation, you can use questions to help guide them. I have found this is a far more effective method than just "telling them" and it need only take 15 minutes if all the solutions and equipment are prepared before hand AND you don't expect them to write it up as a prac (a demonstration takes even less time, but may not be quite as effective, that is, some students will believe you have somehow "tricked" them).
6. Have the students perform the calculations that will enable them to draw a strong monoprotic acid - strong base titration curve. If you have a class of 20 students, they only need to do one of the calculations each, you can tabulate the results and then they can graph the class results. There are a number of reasons for this, it reinforces the nature of the neutralisation reaction, stoichiometry, and of "limiting reagents" and "reactants in excess". It is also enables them to come to a greater understanding of the shape of the curve than if you just present it to them and discuss key points. Finally, if the students do not have a good grasp of why titration curves are the shapes they are, they will have a much harder time coming to terms with the nature of different indicators and why some indicators are more appropriate than others for particular titrations.
7. Do use "real-world" examples. The acetic acid concentration of brands of vinegar is not hard to do, and empowers them (if you have mothers whinging that their daughter will now only let them buy brand X because its better value because it has a higher concentration of acetic acid than other brands, then pat yourself on the back for a job well done!) If you are in a position to be able to safely determine the concentration of sulfuric acid in a lead-acid battery, then this is also not hard to do (but check whether it can be done at your school). Similarly, you will find concentrated HCl(aq) available at you local hardware store (for cleaning bricks) or pool suppliers (for addition to pools) and, if your safety guidelines allow, you can determine the concentrations of these.If you are prepared to take your students through back titrations (indirect titrations) then a wealth of new "real-world" opportunities is open to you.
8. Finally, do not deceive yourself. It is NEVER about what you "teach", it is ALWAYS about what the students "learn". YOU can make up time by giving the students notes you have prepared for them, making them read stuff for homework, making them watch a 30 minute video instead of doing a 2 or 3 day prac, then you can happily tick this off on your list of things to teach, BUT, you must also find out what the students have learned, because you may very well find out that you have been a bit hasty in ticking something off your list! 

Equilibrium Constants and Balanced Chemical Equations

Consider the reaction in which carbon monoxide gas reacts with oxygen gas to produce carbon dioxide gas:

CO_(g) + ½O_2(g) ⇔ CO_2(g)   with equilibrium constant = K₍₁₎ (at temperature = T)

and this reaction:
2CO_(g) + O_2(g) ⇔ 2CO_2(g)   with equilibrium constant = K₍₂₎ (also at temperature = T)

 Is the value of K₍₁₎ the same as the value of K₍₂₎ ?
 Not sure?
That's why AUS-e-TUTE has just added a new set of resources to help you understand this concept!

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 If you are not an AUS-e-TUTE member, there is a "free-to-view" tutorial available for evaluation purposes at http://www.ausetute.com.au/kstoichio.html

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Coral Reef Dissolution

Coral reefs dominate much of the world's tropical coastline, covering about 15% of the seabed shallower than 30 metres.

The largest coral reef in the world is the Great Barrier Reef off the north-east coast of Australia.

What impact will increasing amounts of atmospheric carbon dioxide have on our coral reefs?

Find out in this edition of AUS-e-NEWS!

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Chemical Equilibrium

To understand chemical equilibrium you need to realise that

1. chemical reactions that are in equilibrium DO NOT go to completion. 
2. "reactants" and "products" are all present at equilibrium
3. the "things" that are in equilibrium are the "reactants" and "products" whose concentrations can be changed (that is, substances that exist in solution).
4. equilibrium is the state in which the rate of the forward reaction (reactants → products) is the same as the reverse reaction (products → reactants

Knowing this, we can then write:

• a mass-action expression (http://www.ausetute.com.au/massaction.html)

And then

• the equilibrium law expression (http://www.ausetute.com.au/equilibriumlaw.html)

Which at equilibrium, results in an expression for the

• equilibrium constant (http://www.ausetute.com.au/equicons.html)

Which can be used to:

• calculate concentration (http://www.ausetute.com.au/ricetable.html)
• determine the extent of a reaction (http://www.ausetute.com.au/kextent.html)
• determine the equilibrium constant for the reverse reaction (http://www.ausetute.com.au/kreverse.html)
• determine whether a chemical reaction is exothermic or endothermic (http://www.ausetute.com.au/ktemperature.html)

I.C.E., R.I.C.E and Equilibrium

Need to calculate an equilibrium constant, but, you haven't been given all the equilibrium concentrations?
Then you need an I.C.E. Table (also known as a R.I.C.E. Table, an ICE chart, a RICE chart), and you can find out all about that at AUS-e-TUTE's new tutorial:
http://www.ausetute.com.au/ricetable.html

Now, if you are an AUS-e-TUTE Member, not only can go to the great member's only tutorial, but you can also:

• play the R.I.C.E. game
• answer the R.I.C.E. test questions and get immediate feedback
• do the R.I.C.E. exam
• make a R.I.C.E. worksheet (if you are a Teacher Member)
• do a R.I.C.E. quiz (if you a student with a Class or School Group)

And if you still haven't joined AUS-e-TUTE, you can sign up right here !

Effect of Temperature on Solubility

What happens to the solubility of a salt or a gas at constant pressure when the temperature of the solution changes?
This can be treated as just a special case of Le Chatelier's Principle .....

And AUS-e-TUTe has just added new resources to help you understand the effect of temperature on solubility.
AUS-e-TUTE Members should log-in to use the new tutorial (includes a drill), game, test and exam.

If you are not an AUS-e-TUTE member, there is non-interactive "free-to-view" tutorial currently available for evaluation purposes at http://www.ausetute.com.au/solubilitylcp.html


Get a better understanding of concepts in chemistry, and improve your problem solving ..... join AUS-e-TUTE today!

Calculating the pH of Sulfuric Acid

How do you calculate the pH of a strong diprotic acid, that is, how do you calculate the pH of sulfuric acid?
I'm glad you asked!
AUS-e-TUTE has just added a new set of resources designed to help you answer this question.

AUS-e-TUTE Members should log-in and go to the links for the Polyprotic Strong Acids tutorial, game and test, on the Test Centre homepage, or, alternatively, go to your AUS-e-TUTE syllabus study guide and follow the links from that.

Not a member?
There is a "free-to-view" tutorial currently available for evaluation purposes at
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Calculations for Strong Bases

AUS-e-TUTE has just added new resources!
As part of the update of our "Acids and Bases" topic, tutorials, games, tests, and exams have been added for the following:

• pOH Concepts
• Strength of Bases
• Calculating the pOH of Strong Bases (aqueous solutions)
• Calculating the Hydroxide Ion Concentration of Strong Bases (aqueous solutions)
• Calculating the pH of Strong Bases (aqueous solutions)
• Calculating the Hydrogen Ion Concentration of Strong Bases (aqueous solutions)

Neutral pH?

We hear this term a lot, often in advertising. But what does it mean?

From a Chemist's point of view, there are two different concepts involved in this seemingly harmless "neutral pH" expression. These two different concepts are:

• neutral
• pH

Let's take a look at the Chemist's definition of neutral first.
A solution is neutral if the concentration of hydrogen ions, [H⁺], is equal to the concentration of hydroxide ions, [OH^-].
Chemists often use square brackets to denote concentration, the concentration of  hydrogen ions can be written as [H⁺] and the concentration of hydroxide ions can be written as [OH^-].
So, for a neutral solution:

[H⁺] = [OH^-] = neutral solution 

Pure water is an excellent example of a neutral substance.

Some of the water molecules, actually very few of them, dissociate to form hydrogen ions and hydroxide ions:

H2O

H+ + OH-

Every time a water molecule dissociates, it produces one hydrogen ion, H⁺, and one hydroxide ion, OH^-, so that the concentration of hydrogen ions is always the same as the concentration of hydroxide ions.
Therefore, pure water is always neutral!

The pH of a solution is a measure of the hydrogen ion concentration in the solution. pH can be defined as:

pH = -log₁₀[H⁺]

This equation can be used to calculate the pH of our neutral water, but only if we know the concentration of  hydrogen ions in the water.

The concentration of hydrogen ions in water is not constant!

The concentration of hydrogen ions in water depends on the temperature of the water!
The dissociation of water molecules requires energy:

H2O + energy

H+ + OH-

If you put more energy into the system by heating it, then more water molecules dissociate, the concentration of hydrogen ions increases and the concentration of hydroxide ions also increases.
If you take energy away from the system by cooling it, then fewer water molecules dissociate, the concentration of hydrogen ions decreases and the concentration of hydroxide ions also decreases.
If we were to measure the concentration of hydrogen ions in pure water at various temperatures, we would find the following values:

Water temperature	[H+] x 10−7 M	pH
0°C	0.32	7.50
10°C	0.55	7.26
18°C	0.84	7.08
25°C	1.10	6.96
30°C	1.34	6.87
50°C	2.82	6.55
60°C	3.55	6.46
70°C	4.60	6.34
80°C	5.92	6.23
90°C	7.28	6.14
100°C	8.54	6.07

So what is the pH of water?
The pH of water is dependent on the temperature of the water.
Water is neutral for every value of pH because the concentration of hydrogen ions is always equal to the concentration of the hydroxide ions.
We can ONLY talk about the pH of water IF we state the temperature of the water.
For example, we can talk about water having a pH of approximately 7 at 25^oC, or we could say that the pH of water is approximately 6 at 100^oC.
Pure water is always neutral.
Pure water is neutral at 25^oC.
Pure water is neutral at 100^oC.

As Chemistry students, what we can't say is that water has a pH of 7, or that a neutral aqueous solution has a particular pH, unless we state the temperature of the system.

Further Reading:
Definitions of Acids and Bases
pH
Dissociation Constant for Water

Suggested Study Questions:

1. Plot a graph of temperature versus concentration of hydrogen ions in water. Describe the shape of the line, and write a generalization that links hydrogen ion concentration and temperature.
2. Plot a graph of temperature versus pH of water. Describe the shape of the line and write a generalization linking the  temperature of water and its pH.
3. Use your graph to find the pH of water at:
   • 12^oC
   • 22^oC
   • 32^oC
4. Construct a table giving the concentration of hydroxide ions in water at each of the temperatures shown above.
5. Plot a graph of temperature versus concentration of hydroxide ions in water. Describe the shape of the line, and write a generalization linking hydroxide ion concentration in water and temperature.
6. Use your graph to find the concentration of hydroxide ions in water at:
   • 12^oC
   • 22^oC
   • 32^oC
7. Explain why water is neutral at all temperatures.
8. Explain why the pH of water varies with temperature.

Ammonia Production

Ammonia (NH₃) is one of the most important chemicals in the modern world, mostly due to its use in the manufacture of artificial fertilisers. The Haber, or Haber-Bosch process, is used to produce ammonia and is vital to the production of 100 million tons of fertiliser per year, responsible for sustaining one-third of the Earth's population.

Ammonia is generated naturally by plants and certain bacteria, which extract nitrogen from the atmosphere in a process known as nitrogen fixation. Natural nitrogen fixation occurs at ambient temperatures and pressures, but artificial nitrogen fixation via the Haber-Bosch process requires high pressures (150-250 atmospheres) and high temperatures (300-550 degrees Celsius) to produce the vast quantities of ammonia necessary to satisfy global demand.

The key to the Haber-Bosch process is an iron catalyst which encourages the dissociation of N₂ molecules, and provides a platform on which the resulting N atoms can be successively hydrogenated to yield NH, NH₂ and finally NH₃.

Scientists at the University of Cambridge exposed their iron sample to nitrogen ions, in order to readily build up a coverage of nitrogen atoms on the surface (to a density of just over one nitrogen atom per two top-layer iron atoms at the surface). Under uhv conditions, they can utilise Auger Electron Spectroscopy (AES) to quantify the amount of nitrogen on the surface. Then, they expose the sample to 0.6 mbar H₂ gas for a period of several minutes. This pressure is still very low compared with industrial conditions, but it allows the reaction to proceed sufficiently rapidly for them to take meaningful measurements over a timescale of minutes. If they used only uhv pressures of H₂, the reaction would be so slow that it would take hours, during which time contamination would build up on the surface and ruin the experiment.
After an exposure of several minutes, they rapidly evacuate the experimental chamber to return to uhv conditions and use AES to evaluate how much nitrogen is left on the surface, then expose to H₂ again and repeat. By doing this several times, they can measure the drop in surface nitrogen (corresponding to production of ammonia) as a function of time and temperature.

Their results suggest that, under certain conditions, namely when the ammonia pressure is kept low, the hydrogenation steps (from N to NH to NH₂ to NH₃) may actually be the most important.

Journal Reference:

1. Poobalasuntharam Iyngaran, David C. Madden, Stephen J. Jenkins, David A. King. Hydrogenation of N over Fe{111}. Proceedings of the National Academy of Sciences, 2010; DOI: 10.1073/pnas.1006634107

Further Reading
Haber Process
Nitrogen Cycle

Study Questions

1. Write a balanced chemical equation for the production of ammonia from hydrogen and nitrogen gas.
2. Predict the effect of high pressure in the reaction vessel on the yield of ammonia.
3. The Haber Process is an exothermic reaction. Explain what is meant by the term exothermic.
4. Explain what would happen to the yield of ammonia if the reaction vessel were cooled.
5. It is estimated that between 3 and 5% of the world's natural gas production is used in the production of ammonia. What would the natural gas be used for in the is process?
6. The Haber process typically produces an ammonia yield of between 10 and 20%. Describe 4 ways that this yield could be improved.
7. In the article above it is said that measuring the drop in surface nitrogen corresponds to measuring production of ammonia. Explain why this is true.
8. Describe another way you could measure the production of ammonia.