Chemistry 1020—Lecture 12—Notes

Chapter 5 deals with water. It begins with a discussion of our consumption of water, comparing the various uses for water, the consumption by various societies, and the fact that different uses require water of different qualities.

It then discusses water and its properties, and why it is so special to life on earth. Our bodies are about 60% water, and without a source of water to drink, we would live but a few days.

Has anyone read "The Rime of the Ancient Mariner"?

A poem by Samuel Taylor Coleridge, a late 18^th century poet. One of its stanzas goes:

Water, water, everywhere,

And all the boards did shrink;

Water, water, everywhere,

Nor any drop to drink.

Why no water to drink? Because ocean water is too high in salt concentration for us to be able to use.

Well, not only are there major sources of water with too high a mineral concentration, there are other types of contamination we find in natural waters, and processes we must go through in order to get water sufficiently pure of minerals and other contaminants for the more critical uses. All of these processes which all require the input of energy.

Let’s begin by considering the very unusual properties of water. First, consider what happens to it as heat is added or taken away.

Beginning with ice at –40 ^oC, as you add heat, the temperature increases in an almost linear fashion, with a specific heat of 0.50 cal/g-^oC (or 2.1 J/g-^oC). Then the temperature stays constant while the ice melts, requiring 80 cal/g. (This is called the heat of fusion). Then the temperature of the liquid water increases by one degree per calorie (specific heat of 1.0 cal/g-^oC, or 4.184 J/g—this was the means of defining the calorie). At 100 ^oC, the water turns to steam, requiring 540 cal/g (the heat of vaporization). After all is converted to steam, then the temperature increases again, this time with a specific heat of 0.44 cal/g-^oC.

Sample calculation involving specific heats:

How much heat energy is required to heat your water tank (lets say it holds 80 L) from room temperature (25 ^oC) to 140 ^oF (60 ^oC).

A temperature change of 60-25 = 35 ^oC.

q = heat required = m x D t x sp. ht.

q = 80 L x 10³ mL/L x 1.00 g/mL x 35 ^oC x 1.00 cal/g-^oC

= 2.80 x 10⁶ cal

How much would you save if you heated it to 120 ^oF (48.9 ^oC)

A difference of 48.9-25 = 23.9 ^oC.

q = 80 L x 10³ mL/L x 1.00 g/mL x 23.9 ^oC x1.00 cal/g-^oC

= 1.91 x 10⁶ cal.

a savings of (2.80-1.91) x 10⁶ cal = 8.9 x 10⁵ cal.

Lets convert this to Joules:

8.9 x 10⁵ cal x 4.184 J/cal = 3.7 x 10⁶ Joules or 3700 kJ

Since one kilowatt-hour is 3600 kJ, this is an energy savings of about 1 kilowatt-hours.

Water properties are unusual

All of these properties are very unusual. First of all, water has a very high specific heat (resistance to temperature change)when compared to other substances. That’s what makes it good for calorimetry work. That’s what makes it a good heat exchanger. This high specific heat also helps us to maintain a constant body temperature.

Second, the melting and boiling points of water are unusually high. Normally, boiling points increase with molecular weight. Other substances with molecular weights near that of water (CO₂, CH₄, NH₃, O₂, N₂) are all gases at room temperature.

Third, the heat of fusion of water and the heat of vaporization are both unusually high. This indicates that a lot of energy is required to separate the molecules from each other, i.e. that the intermolecular forces in water are very strong.

Fourth, water is unusual in that it expands when it freezes. Therefore ice floats instead of sinking.

To explain these properties, we need to introduce another concept into our discussion of molecular structure. The concept of electronegativity.

Electronegativity

This term refers to the attraction an atom has for a shared pair of electrons in a chemical bond. Table 5.5, page 156, give values of electronegativity for elements in the first three periods.

Note that electronegativity increases as you move from left to right in the periodic chart. It also increases as you move from bottom to top within a family. Therefore, Fr is the least electronegative element, F is the most electronegative element. (Oxygen is the second most electronegative element.

The distribution of shared electrons between atoms forming a covalent bond is related to the difference in electronegativity. When this difference is great enough, the electron is transferred totally and ions are formed.

In NaCl, the difference is 3.0-0.9 = 2.1, large enough for complete transfer.

In Cl₂, the difference is 3.0-3.0 = 0, so the electrons are shared equally.

In gaseous HCl, the difference is 3.0-2.1 = 0.9, leading to an unequal sharing that puts a partial positive charge in the hydrogen end of the molecule and a partial negative charge on the chlorine. This creates what is called a dipole (like a little magnet).

In water, the H-O bond has an electronegativity difference of 3.5-2.1 = 1.4, and even larger charge separation, also creating a dipole. We call such a bond a polar bond. Covalent bonds can range in degree of polarity, from little or none to larger values, and when the difference gets big enough actually ionic bonds are formed.