Inorganic Compounds

Water’s States: Gas, Liquid, and Solid

The orientation of hydrogen bonds as water changes states dictates the properties of water in its gaseous, liquid, and solid forms.

Learning Objectives

Explain the biological significance of ice’s ability to float on water

Key Takeaways

Key Points

  • As water is boiled, kinetic energy causes the hydrogen bonds to break completely and allows water molecules to escape into the air as gas (steam or water vapor).
  • When water freezes, water molecules form a crystalline structure maintained by hydrogen bonding.
  • Solid water, or ice, is less dense than liquid water.
  • Ice is less dense than water because the orientation of hydrogen bonds causes molecules to push farther apart, which lowers the density.
  • For other liquids, solidification when the temperature drops includes the lowering of kinetic energy, which allows molecules to pack more tightly and makes the solid denser than its liquid form.
  • Because ice is less dense than water, it is able to float at the surface of water.

Key Terms

  • density: A measure of the amount of matter contained by a given volume.

Water’s States: Gas, Liquid, and Solid

The formation of hydrogen bonds is an important quality of liquid water that is crucial to life as we know it. As water molecules make hydrogen bonds with each other, water takes on some unique chemical characteristics compared to other liquids, and since living things have a high water content, understanding these chemical features is key to understanding life. In liquid water, hydrogen bonds are constantly formed and broken as the water molecules slide past each other. The breaking of these bonds is caused by the motion (kinetic energy) of the water molecules due to the heat contained in the system. When the heat is raised as water is boiled, the higher kinetic energy of the water molecules causes the hydrogen bonds to break completely and allows water molecules to escape into the air as gas (steam or water vapor). On the other hand, when the temperature of water is reduced and water freezes, the water molecules form a crystalline structure maintained by hydrogen bonding (there is not enough energy to break the hydrogen bonds). This makes ice less dense than liquid water, a phenomenon not seen in the solidification of other liquids.

Phases of matter: See what happens to intermolecular bonds during phase changes in this interactive.

Water’s lower density in its solid form is due to the way hydrogen bonds are oriented as it freezes: the water molecules are pushed farther apart compared to liquid water. With most other liquids, solidification when the temperature drops includes the lowering of kinetic energy between molecules, allowing them to pack even more tightly than in liquid form and giving the solid a greater density than the liquid.

The low density of ice, an anomaly, causes it to float at the surface of liquid water, such as an iceberg or the ice cubes in a glass of water. In lakes and ponds, ice forms on the surface of the water creating an insulating barrier that protects the animals and plant life in the pond from freezing. Without this layer of insulating ice, plants and animals living in the pond would freeze in the solid block of ice and could not survive. The detrimental effect of freezing on living organisms is caused by the expansion of ice relative to liquid water. The ice crystals that form upon freezing rupture the delicate membranes essential for the function of living cells, irreversibly damaging them. Cells can only survive freezing if the water in them is temporarily replaced by another liquid like glycerol.

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Ice Density: Hydrogen bonding makes ice less dense than liquid water. The (a) lattice structure of ice makes it less dense than the freely flowing molecules of liquid water, enabling it to (b) float on water.

pH, Buffers, Acids, and Bases

Acids dissociate into H+ and lower pH, while bases dissociate into OH and raise pH; buffers can absorb these excess ions to maintain pH.

Learning Objectives

Explain the composition of buffer solutions and how they maintain a steady pH

Key Takeaways

Key Points

  • A basic solution will have a pH above 7.0, while an acidic solution will have a pH below 7.0.
  • Buffers are solutions that contain a weak acid and its a conjugate base; as such, they can absorb excess H+ ions or OH ions, thereby maintaining an overall steady pH in the solution.
  • pH is equal to the negative logarithm of the concentration of H+ ions in solution: pH = – log[H+].

Key Terms

  • alkaline: having a pH greater than 7; basic
  • acidic: having a pH less than 7
  • buffer: a solution composed of a weak acid and its conjugate base that can be used to stabilize the pH of a solution

Self-Ionization of Water

Hydrogen ions are spontaneously generated in pure water by the dissociation (ionization) of a small percentage of water molecules into equal numbers of hydrogen (H+) ions and hydroxide (OH) ions. The hydroxide ions remain in solution because of their hydrogen bonds with other water molecules; the hydrogen ions, consisting of naked protons, are immediately attracted to un-ionized water molecules and form hydronium ions (H30+). By convention, scientists refer to hydrogen ions and their concentration as if they were free in this state in liquid water.

[latex]2\text{H}_2\text{O} \leftrightharpoons \text{H}_3\text{O}^+ + \text{OH}^-[/latex]

The concentration of hydrogen ions dissociating from pure water is 1 × 10-7 moles H+ ions per liter of water. The pH is calculated as the negative of the base 10 logarithm of this concentration:

pH = -log[H+]

The negative log of 1 × 10-7 is equal to 7.0, which is also known as neutral pH. Human cells and blood each maintain near-neutral pH.

pH Scale

The pH of a solution indicates its acidity or basicity (alkalinity). The pH scale is an inverse logarithm that ranges from 0 to 14: anything below 7.0 (ranging from 0.0 to 6.9) is acidic, and anything above 7.0 (from 7.1 to 14.0) is basic (or alkaline ). Extremes in pH in either direction from 7.0 are usually considered inhospitable to life. The pH in cells (6.8) and the blood (7.4) are both very close to neutral, whereas the environment in the stomach is highly acidic, with a pH of 1 to 2.

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The pH scale: The pH scale measures the concentration of hydrogen ions (H+) in a solution.

Non-neutral pH readings result from dissolving acids or bases in water. Using the negative logarithm to generate positive integers, high concentrations of hydrogen ions yield a low pH, and low concentrations a high pH.

An acid is a substance that increases the concentration of hydrogen ions (H+) in a solution, usually by dissociating one of its hydrogen atoms. A base provides either hydroxide ions (OH) or other negatively-charged ions that react with hydrogen ions in solution, thereby reducing the concentration of H+ and raising the pH.

Strong Acids and Strong Bases

The stronger the acid, the more readily it donates H+. For example, hydrochloric acid (HCl) is highly acidic and completely dissociates into hydrogen and chloride ions, whereas the acids in tomato juice or vinegar do not completely dissociate and are considered weak acids; conversely, strong bases readily donate OH and/or react with hydrogen ions. Sodium hydroxide (NaOH) and many household cleaners are highly basic and give up OH rapidly when placed in water; the OHions react with H+ in solution, creating new water molecules and lowering the amount of free H+ in the system, thereby raising the overall pH. An example of a weak basic solution is seawater, which has a pH near 8.0, close enough to neutral that well-adapted marine organisms thrive in this alkaline environment.

Buffers

How can organisms whose bodies require a near-neutral pH ingest acidic and basic substances (a human drinking orange juice, for example) and survive? Buffers are the key. Buffers usually consist of a weak acid and its conjugate base; this enables them to readily absorb excess H+ or OH, keeping the system’s pH within a narrow range.

Maintaining a constant blood pH is critical to a person’s well-being. The buffer that maintains the pH of human blood involves carbonic acid (H2CO3), bicarbonate ion (HCO3), and carbon dioxide (CO2). When bicarbonate ions combine with free hydrogen ions and become carbonic acid, hydrogen ions are removed, moderating pH changes. Similarly, excess carbonic acid can be converted into carbon dioxide gas and exhaled through the lungs; this prevents too many free hydrogen ions from building up in the blood and dangerously reducing its pH; likewise, if too much OH is introduced into the system, carbonic acid will combine with it to create bicarbonate, lowering the pH. Without this buffer system, the body’s pH would fluctuate enough to jeopardize survival.

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Buffers in the body: This diagram shows the body’s buffering of blood pH levels: the blue arrows show the process of raising pH as more CO2 is made; the purple arrows indicate the reverse process, lowering pH as more bicarbonate is created.

Antacids, which combat excess stomach acid, are another example of buffers. Many over-the-counter medications work similarly to blood buffers, often with at least one ion (usually carbonate) capable of absorbing hydrogen and moderating pH, bringing relief to those that suffer “heartburn” from stomach acid after eating.

Overview of the Acid-Base Properties of Salt

Some salts, such as ammonium bicarbonate (NH4HCO3), contain cations and anions that can both undergo hydrolysis.

Learning Objectives

Predict the pH of a solution of a salt containing cations and anions, both of which participate in hydrolysis.

Key Takeaways

Key Points

  • Basic salts result from the neutralization of a strong base with a weak acid.
  • Acid salts result from the neutralization of a strong acid with a weak base.
  • For salts in which both cation and anion are capable of hydrolysis, compare Ka and Kb values to determine the solution ‘s resulting pH.

Key Terms

  • neutralization reaction: a reaction between an acid and a base in which water and a salt are formed
  • hydrolysis: a reaction with water in which chemical bonds break
  • salt: in acid-base chemistry, one of the products in a neutralization reaction

Summary of Acidic and Basic Salts

As we have discussed, salts can form acidic or basic solutions if their cations and/or anions are hydrolyzable (able to react in water). Basic salts form from the neutralization of a strong base and a weak acid; for instance, the reaction of sodium hydroxide (a strong base) with acetic acid (a weak acid) will yield water and sodium acetate. Sodium acetate is a basic salt; the acetate ion is capable of deprotonating water, thereby raising the solution’s pH.

Acid salts are the converse of basic salts; they are formed in the neutralization reaction between a strong acid and a weak base. The conjugate acid of the weak base makes the salt acidic. For instance, in the reaction of hydrochloric acid (a strong acid) with ammonia (a weak base), water is formed, along with ammonium chloride. The ammonium ion contains a hydrolyzable proton, which makes it an acid salt.

Salts in Which Both Ions Hydrolyze

The following is a more complicated scenario in which a salt contains a cation and an anion, both of which are capable of participating in hydrolysis. A good example of such a salt is ammonium bicarbonate, NH4HCO3; like all ammonium salts, it is highly soluble, and its dissociation reaction in water is as follows:

[latex]\text{NH}_4\text{CO}_3(s)\rightarrow \text{NH}_4^+(aq)+\text{HCO}_3^-(aq)[/latex]

However, as we have already discussed, the ammonium ion acts as a weak acid in solution, while the bicarbonate ion acts as a weak base. The reactions are as follows:

[latex]\text{NH}_4^+(aq)+\text{H}_2\text{O}(l)\rightleftharpoons\text{H}_3\text{O}^+(aq)+\text{NH}_3(aq)\quad\quad\text{K}_a=5.6\times10^{-10}[/latex]

[latex]\text{HCO}_3^-(aq)+\text{H}_2\text{O}(l)\rightleftharpoons\text{H}_2\text{CO}_3(aq)+\text{OH}^-(aq)\quad\quad\text{K}_b=2.4\times 10^{-8}[/latex]

Because both ions can hydrolyze, will a solution of ammonium bicarbonate be acidic or basic? We can determine the answer by comparing Ka and Kb values for each ion. In this case, the value of Kb for bicarbonate is greater than the value of Ka for ammonium. Therefore, bicarbonate is a slightly more alkaline than ammonium is acidic, and a solution of ammonium bicarbonate in pure water will be slightly basic (pH > 7.0). In summary, when a salt contains two ions that hydrolyze, compare their Ka and Kb values:

  • If Ka > Kb, the solution will be slightly acidic.
  • If Kb > Ka, the solution will be slightly basic.

Hydrolysis of salts: This video examines the hydrolysis of an acid salt, a basic salt, and a salt in which both ions hydrolyze.