Is Water Weird

Contents

Although oxygen and water appear to be simple substances, both are vital for life and have behaviors that are not necessarily easy to predict. However, a deeper dive into their structures and properties often leads to an understanding of their apparent deviation from normalcy. While Lewis structures are very powerful predictive tools, the structure that results from Lewis structure guidelines for oxygen is one of the very few Lewis structures that incorrectly depicts the electronic structure of a substance. Use of a molecular orbital model approach does correctly indicate that oxygen (O₂) should exist as a triplet state. For water, the Lewis structure approach does correctly result in a model with a bent structure. Consideration of electronegativities of oxygen and hydrogen adds the information that water should be very polar. Using this information coupled with the added feature of hydrogen bonding, it is possible to explain the commonly claimed anomalous behaviors of water.

The models of water below are consistent with the Lewis structure for water.

If hybridization is not considered, the hydrogens should be bonded to the oxygen via p orbitals and the bond angle would be 90^o. For H₂S, this model is very good as the experimental bond angle is 92.1^obut is 14^o off for water. Because there are four groups of electrons around the oxygen, hybridization and electron repulsion models predict that the oxygen should be sp³hybridized and have 109.5 degree bond angles. This model still has a 5 degree discrepancy. Some might argue that the oxygen is only partially hybridized while other arguments for the discrepancy usually suggest that since the electron pairs on the oxygen are not involved in bonding, there is more repulsion than for bonding electrons. The non-bonded orbitals are forced a little farther apart causing the hydrogens to be pushed together. For our purposes, the most important

observation is that the molecule is bent with bond dipoles having the negative terminus on the oxygen starting from each of the hydrogens. Part of these dipoles cancel but there is still a big net dipole directed from midway between the hydrogens towards the oxygen.

Molecular mass. The molecular mass of water, 18 g/mole, is very low for a compound. While there are several elements with low atomic mass values that are solids at room temperature, most compounds with low molecular mass are gases at room temperature. Water being a liquid at room temperature makes water unusual and states of matter for H₂O will be discussed again later. The low molecular mass value impacts many of the properties of water sometimes leading to the claim that water behaves uniquely.

Molarity of Pure Liquid. When we think about the range of possible molarity values, about the highest we commonly encounter in the lab is 18 M for concentrated (98%) sulfuric acid. 47% NaOH also is close to 18 M and 49% HF is 29 M with 90% formic acid also high at 24 M. The values for HF and formic acid are high for the same reason that pure water has the very high molarity of about 56 M. Using the density of water as 1 g/cm³means that the molarity of water can be approximated by the following: (1000g/L)(1 mole/18 g) = 56 moles/L. The very high result is due to division by the low molecular mass value. For students doing problems involving the calculation of molarity, any time a value above 20 M results, the calculation should be rechecked as every answer should be asked if it makes sense.

States of Matter. Ask your self how many compounds you have seen with two states of matter at least temporarily coexisting? If you have ever taken a melting point, both the solid and liquid were present. However, except for water, you probably have not observed substantial amounts of both solid and liquid present at the same time. This is quite phenomenal as there are millions of characterized compounds but only for one of them, namely water, have you seen liquid and solid coexist. Some might claim that their cooking oil sometimes has solid in the bottle of liquid but cooking oil is a mixture and not a single substance. If you have ever soldered a wire connection, the solder was melted but again, solder is a mixture and the liquid quickly solidifies. Since water has a large liquid temperature range (0 – 100^oC) over a mild range of temperatures, it is especially well suited for the establishment of living systems. Hotter temperatures often result in the decomposition of organic compounds (as in cooking) and colder slows reactions down although this does not prohibit life forming reactions. However, the importance of the convenience of the liquid temperature range cannot be understated.

Boiling Point of Water. Water is often compared to compounds of molecular mass like such as methane (CH₄), ammonia (NH₃) and hydrogen fluoride (HF). These three are gases at room temperature. Seldom discussed, it should be noted that LiH and BeH₂ are solids (BH₃ is not included in this discussion as it dimerizes to B₂H₆) with even lower molecular mass values than water. However, LiH is best considered to have ionic bonding compared to the covalent bonding in methane, ammonia, water and hydrogen fluoride. BeH₂ does have covalent bonding but has a two center, three electron bond making it rather unique and not a good model for comparison.

The graph on the above right provides several features that enable meaningful conclusions about water. Delaying temporarily a discussion of the values of the boiling points of water, HF and NH₃, it is observed that the hydrogen compounds of the elements going down each of the groups gradually demonstrate an increase in boiling point as is expected for increases in molecular mass and a corresponding increase in the number of electrons and van-der-Waals forces. Only for methane does the trend continue going to the left to lower mass as water, HF and ammonia have higher boiling points than should be expected based on the hydrogen compounds further down their groups. Clearly, water, HF and ammonia must have a property that causes their significant deviation in trends. All three have polar covalent bonding but the effect is larger than can be accounted for based on common polarity factors. When hydrogen is bonded to F, O or N, a special very strong attraction called hydrogen bonding is possible. Very roughly hydrogen bonds are about 10% of the strength of covalent bonds but significantly stronger than most other types of intermolecular bonds. Hydrogen bonding plays an extremely important role in chemical interactions including a very significant role in the determination of genetics. Hydrogen bonds connect the two strands of the double helix of DNA and determine the pairing of nucleobases that contain the information for repriclation of DNA.

For water, HF and ammonia, because of hydrogen bonding, much more energy (and a higher temperature) is required to break the intermolecular bonds resulting in boiling points much higher than expectations based on molecular mass and polar bonds. If the descending line from H₂Te, H₂Se, H₂S is extrapolated to the left, a predicted boiling point for water results of about -80^o. Hydrogen bonding in water is responsible for a 180^o increase in the boiling point. The boiling points of HF and NH₃ are also much higher than expected.

A comparison of ammonia, water and HF reveals some inconsistencies. The hydrogen bonding strength going across this series should increase as the electronegativity increases the order, N, O, F. However, this is complicated by the number of hydrogens. For water with two hydrogens, more hydrogen bonding with a 3D structure is possible compared to the linear HF resulting in water having a higher boiling point than HF. For ammonia with three hydrogens, the weaker strength of the hydrogen bonding or a less appropriate geometry apparently is more important than an increase in the number of bonds.

The importance of the liquid temperature range of 0 – 100^oC for water cannot be understated. A glance at the liquid temperature ranges of other liquids (see: http://murov.info/orgsolvents.htm) reveals that most other substances with suitable liquid ranges probably would not be too hospitable for life.

Density of Water as a Function of State of Matter. The discussion of hydrogen bonding now leads directly to a consideration of the densities of liquid and solid water. The ratio of the two density values is perhaps the most unusual and difficult to predict property of water. While we observe ice and liquid water mixtures almost every day, few stop to notice the unexpected and very unusual behavior. As mentioned earlier, water is probably the only substance out of millions that you have observed with both liquid and solid states present so that you could notice if the solid sinks or floats in its own liquid. Because it is the only example encountered, most accept the situation as the norm and do not question the very discrepant event. Why was the term discrepant used? Any observation that differs from expectations should be questioned. Often discrepant observations lead to useful new concepts and explanations and should not be overlooked. And yet, the unusual behavior of the ice water system is not noticed by almost everyone.

Starting from scratch, ask if you would expect the solid of a substance to float or sink in its own liquid. As liquids cool, it is expected and almost always true that the volume contracts. Since density is mass divided by volume, a decrease in volume should lead to an increase in density. This trend should continue as the substance freezes resulting in a solid with a higher density than the liquid and the solid should sink. For millions of substances, this prediction is realized. Water is one of the very few substances that behaves contrary to predictions. In the image of 6 test tubes, which one contains water? Only the 3^rd one has the solid floating in its liquid and must be water. The density of water does increase as expected until 4^oC is reached. At this temperature, water begins a very slow decrease in density with temperature until there is almost a 10% expansion when it freezes.

As can be observed in the images above, water molecules in ice have organized into a 3D structure that results in the water molecules being farther apart than they are in liquid water. It is interesting to speculate if life would be the same if water behaved as expected. Where would freezing take place in fresh water lakes in the winter and would life be possible for this situation?

As discussed above, the use of good observational skills often leads to important discoveries. The quotes below from an online PowerPoint presentation http://murov.info/OpticalIllusionsandObservations.ppt emphasize the value of being prepared and careful with observation.

In addition, there are quotations by Sherock Holmes (authored by Arthur Conan Doyle) that are also relevant.

The world is full of obvious things which nobody by any chance ever observes.”

There are 5 elements (arsenic, bismuth, gallium, germanium and silicon) that exhibit the liquid-solid density reversal with only gallium having a melting point near room temperature. There are undoubtedly compounds in addition to water that have this unexpected property, but it is difficult to locate data tables containing liquid and solid densities. https://van.physics.illinois.edu/qa/listing.php?id=1588&t=densities-of-solids-and-liquids.

Specific Heat. Another property of water that is commonly cited as being very unusual is the specific heat of water. It takes 1 cal (or 4.18 joules) to heat one gram of water 1^oC. As you can observe in the table, this value is relatively high compared to other substances.

Like the density reversal, the high value of the specific heat of water affords some advantages. The same amount of energy transferred to 100 g of iron will heat the iron over nine times as many degrees as when transferred to 100 g of water. In effect, this means that water is more resistant to temperature change than most materials. Is it possible to explain the high value for water?

The Dulong and Petit empirical relationship for the product of atomic mass (Mm) and specific heat, (Mm)(Cp) = 25, indicates that the specific heat is inversely proportional to the atomic mass or proportional to the number of moles of particles as with colligative properties. The Dulong and Petit relationship is supported by experimental data as demonstrated in the table below. This means that a measurement of specific heat for an unknown element can help identify the element.

For use of the equation with molecules rather than atoms, the number 25 should be multiplied by the number of atoms per molecule.

a. Is the value 4.184 joules/g deg. for water consistent with this oversimplified approach (note that there is a threshold temperature below which this approach will have significant error)?

b. Compared to metals, what two factors cause the specific heat of water to appear to be exceptionally high?

As water has 3 atoms the constant becomes 3 x 25 = 75. The molecular mass of water calculated from the Dulong and Petit relationship would be 75/4.184 = 17.9 g/mol. This is remarkably close to the actual value of 18.0 g/mol providing support for the modified Dulong and Petit relationship. The specific heat of water is larger than for other chemicals because of the low molecular mass of water and the 3 atoms/molecule and apparently should not be considered discrepant. While the observation of the large value might initially give the impression of weirdness, the value is consistent with expectations.

c. Does this concept work on other substances such as sodium chloride, potassium chloride, calcium chloride, aluminum chloride, carbon tetrachloride or methanol (CH₃OH)? Explain your answer.

Water as a Solvent. Water has several more important life affecting properties such as surface tension that appear to make water distinctive. However, this discussion will not cover other properties but will conclude with a discussion of water as a solvent. Many discussions of water describe it as a universal solvent. That would be extremely weird as it implies that water will dissolve polar and non-polar molecules as well as small and large ones. The statement that water is a universal solvent is a huge misnomer. Since water is a small and very polar solvent, small polar molecules tend to dissolve in water. Solubility is very difficult to predict with the statement “like dissolves like” about as good as you can get. While a good solvent for many ionic compounds, there are many such as AgCl and CaCO₃ that do not dissolve substantially in water. Water is generally not a good solvent for organic compounds that contain five or more carbons and it is a poor solvent for most non-polar organic compounds. Of the millions of known compounds, most are organic and will not dissolve significantly in water. The claim of “universal solvent” to have any meaning must be put into context. There are probably many solvents that dissolve more compounds than water does such as ethanol.

Conclusions. In summary, on the surface, many of the properties of water appear to be distinctive and perhaps unexpected. The answer to the question “is water weird?” depends on perspective. When properties such as a very low molecular mass are taken into account, some of the mysteries can be explained. Perhaps the property that would be the hardest to predict is the reversal of the liquid to solid density ratio. Unless you have some insight into the structures of the solid and liquid phases that result from optimization of hydrogen bonding, this observation would generally not be predicted. Otherwise, the unusual properties such as specific heat are consistent with basic principles of chemistry. If WEIRD means different or unique, then, yes, water is weird. But if weird means unpredictable, then water is not very weird. However, The apparently unusual appearance of the properties gives rise to questions and helps to make the field of chemistry intriguing and interesting.

For web sites that discusses other compounds with unusual properties, please visit: