pH is a term that is placed on a lot of product labels, but most people really don't know what it means.   They usually figure it has something to do with how acidic something is, which is correct.   "pH Balanced" is another term that is thrown about, but also poorly misunderstood.  

Sometimes you see pH values as fractions.   For example, in this scale, baking soda has a pH of 8.3.   If asked what is the H+ concentration of pH 8.3, you would start off the same.  Make it a -8.3 and then make that the exponent of 10.   So 10^-8.3 moles of H⁺ ions per liter is the concentration; however, writing 10^-8.3 is not an acceptable way to report the concentration. Here you must use a calculator and have it calculate 10 with the exponent of -8.3.   That comes out to 5.01x10^-9 moles of H⁺ per liter.   Remember pure water is 1.0 x 10^-7 moles of H⁺ per liter.  So a baking soda solution has about 100 times less H⁺ ions than pure water. This is expected for something that makes an alkaline solution where the OH- concentration is higher than the H+ concentration.

Now let's try tomatoes that has a pH of 4.5.   So that's 10^-4.5 moles of H⁺ ions per liter.  That value should be between 1.0x10^-4 and 1.0x10^-5.  The calculator shows it to be 3.16 x 10^-5, which is between those values. 3.16 x 10^-5 moles of H⁺ ions per liter is about 100 times more than that of pure water. That is expected of anything acidic, which has a pH below 7.

Above when we changed 10^-3 moles per liter into pH 3, the formula we could have used is this:

pH = - log[H⁺]

where [H⁺] is the concentration of H⁺ ion (moles per liter)  The brackets in chemistry usually mean we are talking about concentration.

To change 10^-3 moles per liter into pH 3, let's use the same formula.

pH = - log(10^-3)

The "log of a number" is the exponent of 10 that makes that number.   In the case of 10^-3, it is easy because we see 10 raised to an exponent of -3.  So the exponent is -3.   The formula then says to take the negative of that exponent.  So that would be -(-3), which is just 3.  

Notice the brackets around [H⁺] mean concentration, but when taking the log of a number, parentheses are used instead.  That's why the brackets got changed to parentheses when we wrote pH=-log(10^-3).   The math becomes more difficult if the concentration is not simply a 10 raised to some exponent.   For example, if the concentration of H⁺ ions is given as 3.5 x 10^-4, we will need the entire number (3.5 x 10^-4) written as an exponent of 10.  That's what log function will do.   So below is the formula:

pH = -log(3.5 x 10^-4)

In a spreadsheet the log can be calculated with

=log(3.5*10^-4)

For some calculators you just write log(3.5*10^-4)

For another example, let's see how pH 7 comes from an H⁺ concentration of 1x10^-7M (moles per liter). Using the formula above, it says pH = - log[H⁺]. The [H⁺] is the concentration. So we have pH = -log (10^-7M). To find the log of a number, you find the exponent of 10 that gives you that number.    In this example, a concentration of 10^-7 makes it easy because it is written with 10 with the exponent of "-7". So the log of 10^-7 is simply -7. The formula for pH says pH=-log(10^-7), which becomes pH=–(–7).  So pH = 7.

Why not have pH simply be the concentration and not get complicated using log? The reason is that the concentration of the H⁺ ions in various solutions varies greatly. For example, the concentration of H⁺ ions in your stomach is about billion times more than the concentration of H⁺ ions in an anti-acid remedy like Milk of Magnesia. Instead of working with these large differences in concentration, the log of the concentration makes the values a lot smaller. In other words, pH 1 for stomach acid compared to pH 10 of Milk of Magnesia is easier to work with than using 1x10^-1 (or 0.1) molar H⁺ for stomach acid and 1x10^-10 (or 0.0000000001) molar H⁺for Milk of Magnesia.

In an acidic solution, [H⁺]>[OH^-], therefore the pH is less than 7. In a basic solution, [H⁺]<[OH^-], and the pH is greater than 7. The higher the pH, the more basic the solution: the lower the pH the more acidic the solution.

LITMUS PAPERS: Litmus paper is usually made from about 10 to 15 dyes that are extracted from lichen (a fungus). In the desert we often see rocks with lichen on the shady side of the rock. It's usually green but can also be other colors. The dyes turn blue if in alkaline solutions. However, when in acidic solutions, the dye turns to a red color. So blue litmus paper turns red if the solution is moderately to highly acidic (pH<4.5). Red litmus paper turns blue if in an alkaline solution (pH>8.3). In between these pH levels it is purple.

The word "Litmus" is from an Norwegian word, "Litmos".  "Lit" means "light" or "colored" and "mos" means "moss", a plant similar to lichen.

Universal pH Paper: To the left is a typical universal pH paper that measures the "full" pH range of "1 to 14" or "0 to 14". This paper usually has 4 squares to indicate the color. Each square has a different dye because different dyes will change colors over different ranges of pH. You match the 4 colors with the closest 4 colors on the pH chart that comes with the pH paper.

BLUEBERRY EXTRACT: To the left are colors of blueberry extract (blueberry juice) at different pH levels.  Notice that natural blueberry juice has a pH of about 5.3 (slightly acidic). At more acidic levels is it more red.   At alkaline pH it is more brown.    As you can see the color change isn't as dramatic as it is for red cabbage extracts.

Extracts from red grapes, apple skins, other colored fruit and vegetables will change their colors depending on the pH.   Making an extract can be done by chopping or grinding the plant, then boiling it in water or soaking it in a solvent like acetone.  Either way, this extracts the colored pigments.

The old favorite poem starter "Roses are Red, Violets are Blue..." has a new explanation. The pH in violets is probably different than the pH in roses.

There's been an effort to genetically modify roses to make them blue, but they haven't yet controlled the pH, so no success yet.

If the colored pigments in flowers were extracted, many of those would probably be good as pH indicators. 

Wikipedia lists about 25 man-made chemicals that change color with pH (pH indicators).     Here is the URL if you want to check it out:

http://en.wikipedia.org/wiki/PH_indicator

Bromophenol Blue:    This pH indicator is in your kit.  It was used in lab 7 to check the acidity of water that reacted with sulfur dioxide that came from burning sulfur. It caused the water to turn yellow which means the pH was at 3.0 or lower (more acidic). Notice in the image the yellow color at pH 3.0 or lower.

The name "phenol" refers to the rings that have an "OH" group on it.   Phenol is a solvent that has one 6-carbon ring with a single OH group attached.

You may notice that like the pH indicators from plants, these man-made pH indicators also have 6-carbon rings that have double bonds alternating with single bonds.   Visible light can cause these electrons to jump to higher energy levels, which means the color that causes that jump gets absorbed.  So the reflected color is no longer white.  For compounds that don't have these alternating double bonds, only high energy ultraviolet light can cause the electrons to jump to higher energy states.   Visible light is not absorbed, so compounds without the alternating double bonds appear clear or white.  Groups like OH, COOH, CO, and even nitrogen atoms are affected by H+ and OH- ions in solution because H+ and OH- ions may add or remove charges on the indicator molecule.  Different charges on the molecule causes it to absorb different frequencies of light, so it changes color.

Bromothymol Blue:   This pH indicator is in your kit.   Because it changes colors in slightly acidic to slightly alkaline solutions, it is used for situations that are checking pH places like swimming pools or fish tanks where the pH shouldn't vary too far from neutral pH (7).   

Looking at the photo, notice that at pH 6.0 and below a solution with bromothymol is yellow.  At a neutral pH it is green. At 7.6 and above it is blue.  So it changes color over a rather narrow range of H+ concentration.  In this lab we will show that the CO2 in your breath can change slightly alkaline water into acidic water.

Bromothymol blue is sometimes used prior to a birth to detect if the membrane surrounding the baby has broken.  The amniotic fluid that surrounds the baby has a pH above 7.6.   So it would turn the Bromothymol blue to a blue color.  

The molecular structure of Bromothymol Blue is similar to bromophenol above except two bromine atoms were replaced with an isopropyl group (a carbon with two methyl (CH₃) groups attached.  Isopropyl groups look like "Y's" in the image.

"Thymol" is in the name "bromothymol" because thymol has a 6-carbon ring with bromine, OH, and isopropyl groups.   As the name suggests, thymol comes from the spice called "thyme".   Thymol, by the way, has antibacterial qualities.

The below image shows color changes even in the narrow range of pH 6.2 to 7.5.

Phenolphthalein:  This pH indicator is also in your kit.    This was used in Lab 4 to show that when sodium bicarbonate or sodium carbonate dissolve, they cause excess OH- (hydroxide) groups to form in the water.  That makes the water alkaline.   The OH- groups also pull off the hydrogen atoms of the OH groups on the phenolphthalein molecule.    That changes the arrangement of double bonds on the phenolphthalein molecule so that they alternate throughout the molecule.  That gives it the ability to absorb all colors except pink, which makes the solution pink (reddish purple=fuchsia/pronounced few-shuh).

Phenolphthalein is used to show when acids get neutralized because the solution will go from clear to pink immediately after the acid gets neutralized.

Using a buret (shown) to drip in an alkaline solution like sodium hydroxide (NaOH), the amount of NaOH used to neutralize the acid can be measured at the point the solution with phenolphthalein goes pink.   That will allow a calculation of how much acid was present.  This is useful for determining the strength of acids like battery acid, vinegar (acetic acid), and even organic acids like Aspirin and citric acid. 

Below is a benchtop pH meter.  Notice this is accurate to the thousandths place. By changing the probes, these meters can read the concentration of other ions other than just H+ ions.

Phosphoric acid:  H₃PO₄

In water, about one out of every 10 phosphoric molecules dissociate into an H⁺ ion and a
H₂PO₄^- ion.

H₃PO₄ → H⁺ + H₂PO₄^-

This means if you had 20 phosphoric molecules, there would be two that would have dissociated into H⁺ and H₂PO₄^- ions.

It usually takes two compounds to do this.  If this is a buffer that is trying to a hold a specific acidic pH (less than pH 7), then it will contain a weak acid that releases H⁺ ions if the number of H⁺ ions go down, and another compound that can absorb H⁺ ions if the number of H⁺ ions increase.  

First let's look at what a weak acid does.  One example is phosphoric acid (H₃PO₄). which is used in sodas to make them taste more tart.  Phosphoric acid is called an acid because the hydrogen atoms can break away from the phosphoric acid molecule while leaving their electrons behind.  So the hydrogen atoms become H⁺ ions.   If one hydrogen leaves, the other part of the phosphoric acid now has a negative charge and is now called an dihydrogen phosphate ion (H₂PO₄^-).

Phosphoric acid is called a weak acid because it doesn't release all of its hydrogen atoms (as H⁺ ions) into the water. Only about 1 out of every 10 phosphoric acid molecules will be dissociated into H₂PO₄^- and H⁺.   Strong acids like hydrochloric acid (HCl) release all of its hydrogen atoms (as H⁺ ions) into the water.

A weak acid is actually a pretty good buffer if it only encounters something alkaline that neutralizes or removes the H⁺ ions in the solutions.   Consider an example of putting 20 molecules of phosphoric acid into water.   Remember, phosphoric acid has the tendency to release one H+ ion for ever 10 molecules of phosphoric acid.  So in this case 20 molecules of phosphoric acid there would be 2 H⁺ ions and 2 H₂PO₄^- ions.

H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄
 H⁺ H₂PO₄^-   H⁺ H₂PO₄^-

You might think that if you added 5 OH^- ions, that the two H+ ions in solution would get turned into water by combining with 2 OH^- ions. Then the pH would go from acidic to alkaline because there would then be 3 OH^- ions in the solution.  Notice below that two H⁺ and OH^- ions have formed water (HOH), which is another way to write H₂O. 

H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄ H₃PO₄  H₃PO₄ H₃PO₄  H₃PO₄ H₃PO₄ H₃PO₄ H₃PO₄  H₃PO₄ H₃PO₄
HOH H₂PO₄^-   HOH  H₂PO₄^- 
OH^- OH^- OH^-

The 3 OH^- ions do not stick around because 3 phosphoric acid molecules can give one H⁺ ion each directly to the 3 OH^- ions turning the OH^- ions into water (HOH).

3H₃PO₄ + 3OH^- → 3H₂PO₄^-   + 3HOH

At this point our solution has 15 phosphoric molecules, 5 dihydrogen phosphate ions, and 5 extra water molecules. 

H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄   H₃PO₄  H₃PO₄  H₃PO₄ H₃PO₄  H₃PO₄  H₃PO₄   H₃PO₄H₃PO₄  H₃PO₄
H₂PO₄^- H₂PO₄^-  H₂PO₄^-   H₂PO₄^-   H₂PO₄^- 
HOH   HOH  HOH  HOH   HOH 

 Notice above that there are no H⁺ ions or OH- ions.  So it seems that the solution would be neutral (pH 7); however, recall that about 1 out 10 phosphoric molecule will dissociate to form H⁺ and H₂PO₄^-- ions.   So with 15 molecules of phosphoric acid left, they will on average keep 1.5 H⁺ ions in solution.  Therefore the concentration went from the original 2 H⁺ ions in solution to 1.5 ions in solution.   So the pH didn't change much.   So that shows that phosphoric acid or other weak acids can buffer against addition of something alkaline like OH-.   However, by itself, weak acids cannot keep the pH constant if acids (more H⁺ ions) are added.   Let's examine that situation.

Below again is what 20 phosphoric acid molecules look like in water.

H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄
 H⁺ H₂PO₄^-   H⁺ H₂PO₄^-

Now if 6 molecules of hydrochloric acid (HCl) is added, that will introduce 6 more H⁺ ions to the 2 H⁺ ions already present.    The only thing that can absorb these extra H⁺ ions are the negative dihydrogen phosphate ions (H₂PO₄^-).  But there are only two of them.  So below is what happens with 6 H+ ions.  Notice that 2 of these 6 turn the H₂PO₄^- into H₃PO₄ (phosphoric acid).  Note, I am writing the newly formed H₃PO₄ (phosphoric acid) as HH₂PO₄, just to show that one of these hydrogen atoms came from hydrochloric acid.  Below is the result.

H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄  H₃PO₄
H⁺  HH₂PO₄  H⁺   HH₂PO₄
H⁺  H⁺  H⁺  H⁺

Now we went from 2 H+ ions to 6 H+.  So the concentration of H⁺ tripled, which means the pH went up significantly, too.   So phosphoric acid by itself is not a good buffer against the addition of acids.   One fix is to have more of the H₂PO₄^- ions that can grab H⁺ ions.  We can supply those H₂PO₄^-  ions with another compound such as sodium dihydrogen phosphate (NaH₂PO₄). This compound dissolves in water to form Na⁺ ions and the  H₂PO₄^- ions we want.  So by having a mixture of phosphoric acid and sodium dihydrogen phosphate (NaH₂PO₄), the phosphoric molecules can replenish H⁺ ions if something alkaline is added, and the extra H₂PO₄^- from NaH₂PO₄ can absorb H⁺ ions if an acid is added.  That is a true buffer.

Buffering actions are very important in the human body because enzymes are only active in a narrow range of pH.   The pH value of blood must be maintained close to 7.4 (slightly basic/alkaline).  Buffers in blood keep the pH value almost constant even if acidic foods are eaten or lactic acid is produced in the muscles during exercise.  The compounds in blood that act as buffers are H₂CO₃ (carbonic acid) and HCO₃^- (bicarbonate ions).   As you might guess, H₂CO₃ can replenish H⁺ ions if something alkaline is introduced and  HCO₃^- can absorb extra acids (H⁺). 

H₂CO₃ → HCO₃^- + H⁺

HCO₃^- + H⁺ → H₂CO₃

In the online CHM107 lab, which is also done at home, fresh red cabbage is chopped up and then cooked on a stove top. The boiling breaks down the cell walls allowing the dye to dissolve in the water.  The water is then poured in a beaker and then passed through a filter.  That liquid extract is then used to test the pH of various items.  Here in this online CHM130 lab, you will be using red cabbage that has been chopped and dried.  So you don't have to go out to buy fresh red cabbage.

Get one of the tea candles and light it.   Use the chrome test tube holder to hold the test tube while the flame heats the bottom of the test tube.

Do not point the opening of the test tube towards your face or any one that might be standing by. The boiling may cause it to splatter out the end of the test tube.

It the heating looks like it will boil over, back away from the heat and let it cool a few seconds before heating it more.

You will be measuring the pH of the following household items, which are all in your kit (even those assembled for Fall 2014).

The pH of each will be measured by using a pH Meter, red cabbage extract, three man-made pH indicators, and Universal pH paper.

Since there are six items to test and six ways to measure pH, that means there will be 36 (6 x 6) tests to do.  The challenge is to stay organized and to record the results as you go.    Typically, these kind of multiple tests are done using porcelain spot plates (see below).  Spot plates allow you to do several tests on the same plate.  Unfortunately, these have only 12 depressions.  So instead of the white porcelain spot plates, your kit uses plastic reaction well plates which have 24 depressions and therefore can test 24 samples (left image).   These also have a lid.  These are also called "reaction well plates", "culture plates", or "microplates".

On the left are the six household products that you will be testing.   The bottom test tube is one labeled lemon juice.   Line them up in the order shown.  That way they will match the order shown on the paper that labels the cups, spot plates, and microscope slides.

Note that each test tube should have its own pipette.    

Do the same thing you did with sodium carbonate with the second test tube (sodium bicarbonate/NaHCO₃).    

1.  Use the microspatula to remove a small amount of it and put it in the cup that is sitting on the circle labeled NaHCO₃.   

2.  Add some purified water to it. (about the same amount as was done for sodium carbonate) 

3.  Use a clean pipette to draw the liquid into the pipette and then squeeze the liquid back out to stir the contents.  Doing that a few times will dissolve the sodium bicarbonate (NaHCO₃). 

4.  Place the pipette back alongside the sodium bicarbonate test tube.  Don't get the pipette mixed up with the ones used for the other products.

The next test tube is labeled "Acetic Acid (Vinegar)".   Pour all of its contents into the cup that is on the line marked "Acetic Acid". 

The last test tube is "Lemon Juice".  Pour all of the lemon juice into the cup that sits on the line marked "Lemon Juice". 

Now with all of these cups filled with their respective household product, you are ready to measure their pH using the pH meter.

Test the first cup, which is the one with the dissolved sodium carbonate (washing soda).  Noticed that the cup was tilted to get the liquid deep enough to cover the glass sensor probe.   In this image, one had was used because the other hand was busy taking the photo.  In your case, use two hands.

After rinsing with tap water, it's a good idea to rinse off the glass probe with some distilled water.  However, if you are low in distilled water, you can skip that.  

The pH meter is now ready to measure the pH of a different liquid.  In this case it will be the sodium bicarbonate.

Now that you've finished measuring the pH of these liquids using a pH meter, you can use the liquids to check pH with the use of pH indicators.

Using the pipette that was used to mix the sodium carbonate, use it to transfer some of the dissolved sodium carbonate (Na₂CO₃) to the 4 spot plate wells on the top row of the spot plate.

Here you can see the first column is titled "Red Cabbage".   A drop or two of the red cabbage extract has been added to each of the wells in the left column.   The color will be an indicator of the pH of each of the samples.   Write the color in the data table. Also, using the chart below, estimate the pH. 

These are the colors that were recorded that described the color the samples turned to when the red cabbage extract was affected by the pH of each sample.   This also shows the estimated pH that the color represents.

Bromothymol Blue: 

Now you will do the same procedure with each of these liquid pH indicators.   Bromothymol Blue is the next column.  So do it next.   Again, one or two drops in each well (spot) is all you need.

These are the colors that the Bromothymol Blue indicator became when affected by the pH of the samples.  This type of indicator is used when you only want to know if the pH is staying mostly neutral, such as in aquariums or swimming pools.  If the water starts to become acidic, the indicator will be yellow.  If the water is becoming alkaline, it will turn blue.  So the blue colors of Na₂CO₃, NaHCO₃ and Milk of Magnesia mean those are all above a pH of 7.6, meaning they are all alkaline.   Milk is around a pH of 7, so it's color is a pale yellow-green (neither yellow or blue).    The acetic acid and lemon juice are definitely yellow and therefore have pH levels below 6.0, which means they are acidic.  Again, it doesn't tell you how acidic, just that it is.

Bromophenol Blue:  Now use the Bromophenol Blue bottle to add one or two drops to each sample for its column.   The image shows what affect the pH of the sample had on the Bromophenol Blue indicator.   For this indicator, it becomes yellow if the pH is at 3.0 or below (more strongly acidic).   We see that acetic acid and lemon juice are yellow.  So those are at or below pH 3.0.   The others are all purple or bluish purple.  So they are greater or equal to pH 4.6.

Record the color of each sample in the appropriate place on the data table sheet and the pH that its color indicates.

Phenolphthalein:  The last liquid pH indicator is phenolphthalein.   It is clear in neutral or acidic solutions but pink or reddish purple (fuchsia) in alkaline solutions above pH 8.2.   Earlier, we found that the solution of NaHCO₃ (sodium bicarbonate) had a pH of 8.4 according to the pH meter.  That means it is slightly more alkaline than the pH 8.2 that phenolphthalein starts to turn color.  That's why it is pink and not a darker purple (fuchsia) like the milk of magnesia and washing soda (Na₂CO₃).  Those had pH levels above 10 (about 100 times more OH- ions than baking soda (NaHCO₃/sodium bicarbonate).

Phenolphthalein is good for quantitative analysis of acids because it goes from clear to pink when it crosses from being acidic to being alkaline with only a fraction of a drop of an alkaline solution.

Record the color for each sample in the appropriate box on the data table sheet and the pH that the color indicates.

Your last measurements of pH will use the Universal pH strips.   Lay one alongside the drop on the microscope slide.  The solution will soak into the pH strip and the strip will change colors unless it is neutral.

Take each strip off of the microscope slide one by one and line it up to the best match on the Color Chart for pH Test Papers.     The colors on the wet strip will be look more saturated than the colors on the chart, but use your best judgment to find a match.

Record the pH on the data table sheet.