Pure Water’s pH is only 7.00 at a Specific Temperature.

Pure water’s pH is only 7.00 at a specific temperature of 25.0 °C.  Students (and teachers too) hear that number so often that they forget where it comes from. And forgetting its origins makes one forget that if the temperature deviates significantly from 25.0 °C, you will get unfamiliar numbers for the pH of pure water.

At any temperature, pure water will always have the same concentration of ions resulting from a very slight splitting of the life-essential molecule into a positive hydrogen ion and a negatively charged hydroxide ion. The product of each ion’s concentration will equal its so-called Kw of 1.01 × 10-14 at 25.0 °C. To calculate the concentration of H+ ions, you first take the square root 1.01 × 10-14  and then take the negative logarithm of H+, the definition of pH. It yields 7.00.

But changing temperature usually affects any equilibrium constant, including Kw. In this case raising temperature helps water split up. You get more ions, thus a higher ion product. Kw becomes 5.48 × 10-14 at 50 ° C, Lowering the temperature has the opposite effect on equilibrium, and Kw becomes 0.29 × 10-14 at 10° C.

When you recalculate pH of pure water at 50 ° C and 10° C, we obtain pH’s of 6.63 and 7.27, respectively. The temperature does not make water either slightly acidic or alkaline. It’s just that the middle or neutrality point of the pH scale at those different temperatures changes. The 7.00 is not set in stone. The middle point of the pH scale is a setting derived from what the Kw happens to be at 25.0° C.


The pH is also 7.00 for aqueous solutions whose solutes at that same temperature do not affect the ions that water itself produces. When the temperature changes for those solutions, the pH will change accordingly. In our bodies, if temperature was the only factor, then out physiological pH would be below 7.00. But the presence of bicarbonate ions eats up some of the hydrogen ions, setting the physiological pH at about 7.4. The  pH of the extracellular fluid of tumour cells, as determined by probing microelectrodes, is acidic. That truth has been known for at least 3 decades, and of course nothing one eats will have any impact on the pH of that fluid.

Often teachers have to reiterate to get ideas across. So indulge me. 🙂 Does temperature affect neutrality of pure water? No. The concentration of hydrogen ions will be equal to that of hydroxide ions as long as no solute interferes with one of them. Does raising temperature raise the concentration of OH? Yes. Of H+? Yes. Will that in turn affect pH? Of course, by definition.


The Periodic Table of the Elements’ Natural Sources

EMA1EtCWsAIopQaThere are thousands of different periodic tables in existence. Aside from the usual ones that offer atomic masses and numbers, for a long time we have had those that revealed various periodic trends. In more recent years, some have focused on their time or place of discovery, on cosmic origins of the elements, and even on endangered “species”.

Many academic institutions and a slightly-richer-than-the average-guy by the name of Bill Gates,  have placed actual samples of elements in cubicles to create a 3D-version. There are more modest tables filled with beautiful photographs of the  elements—in fact, you can even get a 1000 piece jigsaw puzzle version.

(Although it should have had 118 or 1180 pieces to match the number of elements. 🙂 ) 914eg7tsctl

I thought of creating one more, a table that focuses on some of the natural sources of the elements—even though I’m sure the idea is far from original. You probably know already that such a table will leave out synthetic ones, about 34 of them. Of course, for most of the remainder, there is more than one natural source. So anyone else who has created such a table will have put together merely one of at least millions of possibilities.


You will notice that whereas a periodic table has mostly metals, the natural source- version consists of mostly minerals, which I find more aesthetically pleasing. A rock is a heterogeneous mixture of minerals, while a mineral is similar to a chemical compound, but it is not as narrowly defined. Its composition can vary within limits; impurities can drastically change the colour of a mineral, and those impurities can sometimes be the only source of the element, as is the case with rhodium and several of the rare earth elements.

The list of elements that can be found in their native, non-bonded state is longer than most of us imagine. It includes four of the five elements in group 15: nitrogen, arsenic, antimony and bismuth; all three mintage elements: copper, silver and gold; iron and nickel in meteorites—in fact native iron can also be found in basalt; five other heavy metals: osmium, rhodium, iridium, palladium and platinum; oxygen, sulfur and, of course, all six noble gases.

For four of the elements, helium, gallium, rubidium and cesium,  I included spectra, which is how those elements were discovered. In helium’s case, the scientists were looking at the sun’s outer layers during the eclipse of 1868. Only decades later was helium gas found on earth when it was found to be released from a uranium ore. Soon after, they realized that lots could be extracted from natural gas sources.

If you refer back to the first table I listed, that of the endangered elements, you will notice that helium is one of them. The number of suppliers worldwide is limited. If one  of them experiences issues, shortages quickly develop. This leads to a spike in prices for the simplest but most essential of the noble gases. Helium is used as a coolant in MRIs, smartphone-manufacturing and other applications.

It’s just not recycled enough, if at all, as is the case with many of the endangered elements.

Thirteen (mostly) Chemistry Demonstrations in 280 Characters Or Less

  1. Ground #helium balloons with a bunch of grapes. Then remove 1 grape at a time until the buoyant & gravitational forces balance out. The balloons will be suspended in the air. Children quickly catch on and have injected a little #science into an otherwise dull wedding reception.
  2. Add copper to HCl & watch nothing happen. Add Cu to HNO3, & NO2 or NO forms, depending on the acid’s concentration. Add CuO to citric acid, wait a few days & a (patina-like?) material forms.
  3. ZNO44To a beaker, add sand and a 50% solution of methanol. Then add spatula tips of zinc oxide powder. Close lights. Transitions are very temperature-dependent and different parts of the flame create a variety of colors. Students prefer demo to drugs.
  4. Add calcium to water & phenolphthalein. Collect H2. Ignite it. Color change in solution reveals hydroxide formation. White precipitate of CaO settles below fuchsia solution. Filter it. Blow into solution of Ca(OH)2 to form CaCO3. More CO2 forms acid, gets rids of cloudiness.
  5. ammoniaAdd 2 drops of bromothymol blue to a (pH ~ 4) solution in a flat-bottomed flask. Add dilute NaOH to beaker. Bring the flask to a boil for 3-5 minutes. Remove heat source & wrap a cold, wet rag around the flask. Be awed by work of ΔPV.
  6. Get a hand-held digital microscope. Use it to reveal  the sensuous surface of a grapefruit, an aborted seed; and the oxidation of copper in an old penny. 

SANYO DIGITAL CAMERAChip off 5 samples from a boulder. Use water displacement in large cylinder to find the volume of each piece. Mass each rock. Obtain average density. With latter & an estimate of the boulder’s volume, get an estimate of the boulder’s mass.
8. Get a thin flow from a water tap. Wrap cotton shirt around a plastic comb. Rub it. Move the comb towards the water without touching it. Watch the stream bend like a banana. Water is neutral, but something charged within it is attracted to the oppositely charged comb.Static 9. Demonstrate that old pieces of magnesium often won’t flash in a Bunsen burner flame. Their surface has reacted with air. Wipe a piece with a paper towel that’s wet with dilute acid. Dry, weigh, ignite& look away! After it flashes, reveal that white residue’s mass > than original.
10. CO2Add two drops of bromothymol blue to 4 different test tubes containing tap water. Add distilled water to the 1st; it remains green. Add baking soda to 2nd, get blue. To the 3rd and 4th add vinegar & dry ice (CO2). Both go yellow as both additives lead to H+.

11. Ignite a hydrogen-filled balloon. Note red color from excess H2 incandescing in heat of reaction. Fill a 2nd balloon with 2:1 ratio of H2 to O2 & ignite; observe no red color. Fill 3rd balloon with hydrogen and add a little copper sulfate. Explosion becomes green-colored.

12. Spread a few grams of iron filings on a filter paper. Place it a magnetic stirrer. Turn it on at medium speed. A beautiful display of a magnetic field in motion ensues. It resembles a living colony of microorganisms.

13. The calcium carbonate in blackboard chalk is long-baked, so it quickly settles, and more stays out of the respiratory system. Yet, if you examine any ledge or border high above the board, it fills with chalk dust over the school year. How? Brownian motion.

14. Imagine someone you has wronged you. Imagine taking a wet sheet of newspaper& sticking it on the guy’s windshield on a cold day. Due to H-bonding he’ll never be able to scrape it off, unless he has access to hot water. Imagine writing on the paper, “Revenge is best served cold.”

More to come.


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