## 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.

$2\text{H}_2\text{O}\leftrightharpoons\text{H}_3\text{O}^{+}+\text{OH}^-$

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.

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.

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.

## Water’s Cohesive and Adhesive Properties

Cohesion allows substances to withstand rupture when placed under stress while adhesion is the attraction between water and other molecules.

### Learning Objectives

Describe the cohesive and adhesive properties of water.

### Key Takeaways

#### Key Points

• Cohesion holds hydrogen bonds together to create surface tension on water.
• Since water is attracted to other molecules, adhesive forces pull the water toward other molecules.
• Water is transported in plants through both cohesive and adhesive forces; these forces pull water and the dissolved minerals from the roots to the leaves and other parts of the plant.

#### Key Terms

• adhesion: The ability of a substance to stick to an unlike substance; attraction between unlike molecules
• cohesion: Various intermolecular forces that hold solids and liquids together; attraction between like molecules

Have you ever filled a glass of water to the very top and then slowly added a few more drops? Before it overflows, the water forms a dome-like shape above the rim of the glass. This water can stay above the glass because of the property of cohesion. In cohesion, water molecules are attracted to each other (because of hydrogen bonding), keeping the molecules together at the liquid-gas (water-air) interface, although there is no more room in the glass.

Surface Tension: The weight of the needle is pulling the surface downward; at the same time, the surface tension is pulling it up, suspending it on the surface of the water and keeping it from sinking. Notice the indentation in the water around the needle.

Cohesion allows for the development of surface tension, the capacity of a substance to withstand being ruptured when placed under tension or stress. This is also why water forms droplets when placed on a dry surface rather than being flattened out by gravity. When a small scrap of paper is placed onto the droplet of water, the paper floats on top of the water droplet even though paper is denser (the mass per unit volume) than the water. Cohesion and surface tension keep the hydrogen bonds of water molecules intact and support the item floating on the top. It’s even possible to “float” a needle on top of a glass of water if it is placed gently without breaking the surface tension.

These cohesive forces are related to water’s property of adhesion, or the attraction between water molecules and other molecules. This attraction is sometimes stronger than water’s cohesive forces, especially when the water is exposed to charged surfaces such as those found on the inside of thin glass tubes known as capillary tubes. Adhesion is observed when water “climbs” up the tube placed in a glass of water: notice that the water appears to be higher on the sides of the tube than in the middle. This is because the water molecules are attracted to the charged glass walls of the capillary more than they are to each other and therefore adhere to it. This type of adhesion is called capillary action.

Adhesion: Capillary action in a glass tube is caused by the adhesive forces exerted by the internal surface of the glass exceeding the cohesive forces between the water molecules themselves.

Why are cohesive and adhesive forces important for life? Cohesive and adhesive forces are important for the transport of water from the roots to the leaves in plants. These forces create a “pull” on the water column. This pull results from the tendency of water molecules being evaporated on the surface of the plant to stay connected to water molecules below them, and so they are pulled along. Plants use this natural phenomenon to help transport water from their roots to their leaves. Without these properties of water, plants would be unable to receive the water and the dissolved minerals they require. In another example, insects such as the water strider use the surface tension of water to stay afloat on the surface layer of water and even mate there.

Cohesion & Adhesion: Water’s cohesive and adhesive properties allow this water strider (Gerris sp.) to stay afloat.

## Water’s High Heat Capacity

Water is able to absorb a high amount of heat before increasing in temperature, allowing humans to maintain body temperature.

### Learning Objectives

Explain the biological significance of water’s high specific heat

### Key Takeaways

#### Key Points

• Water has the highest heat capacity of all liquids.
• Oceans cool slower than the land due to the high heat capacity of water.
• To change the temperature of 1 gram of water by 1 degree Celsius, it takes 1.00 calorie.

#### Key Terms

• heat capacity: The capability of a substance to absorb heat energy
• specific heat: the amount of heat, in calories, needed to raise the temperature of 1 gram of water by 1 degree Celsius

### Examples

The high heat capacity of water has many uses. Commercial nuclear reactors release large amounts of thermal energy (heat) during radioactive decay of fission products. The heat is quickly transferred to a pool of water to cool the reactor. The water then remains hot for a long time due to its high heat capacity.

The capability for a molecule to absorb heat energy is called heat capacity, which can be calculated by the equation shown in the figure. Water’s high heat capacity is a property caused by hydrogen bonding among water molecules. When heat is absorbed, hydrogen bonds are broken and water molecules can move freely. When the temperature of water decreases, the hydrogen bonds are formed and release a considerable amount of energy. Water has the highest specific heat capacity of any liquid. Specific heat is defined as the amount of heat one gram of a substance must absorb or lose to change its temperature by one degree Celsius. For water, this amount is one calorie, or 4.184 Joules. As a result, it takes water a long time to heat and a long time to cool. In fact, the specific heat capacity of water is about five times more than that of sand. This explains why the land cools faster than the sea.

$\displaystyle{C=\frac{Q}{\Delta T}}$

The resistance to sudden temperature changes makes water an excellent habitat, allowing organisms to survive without experiencing wide temperature fluctuation. Furthermore, because many organisms are mainly composed of water, the property of high heat capacity allows highly regulated internal body temperatures. For example, the temperature of your body does not drastically drop to the same temperature as the outside temperature while you are skiing or playing in the snow. Due to its high heat capacity, water is used by warm blooded animals to more evenly disperse heat in their bodies; it acts in a similar manner to a car’s cooling system, transporting heat from warm places to cool places, causing the body to maintain a more even temperature.

## 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.

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.

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.

## Water’s Solvent Properties

Water’s polarity makes it an excellent solvent for other polar molecules and ions.

### Learning Objectives

Explain why some molecules do not dissolve in water.

### Key Takeaways

#### Key Points

• Water dissociates salts by separating the cations and anions and forming new interactions between the water and ions.
• Water dissolves many biomolecules, because they are polar and therefore hydrophilic.

#### Key Terms

• dissociation: The process by which a compound or complex body breaks up into simpler constituents such as atoms or ions, usually reversibly.
• hydration shell: The term given to a solvation shell (a structure composed of a chemical that acts as a solvent and surrounds a solute species) with a water solvent; also referred to as a hydration sphere.

### Examples

• Sugar, sodium chloride, and hydrophilic proteins are all substances that dissolve in water.
• Oils, fats, and certain organic solvents do not dissolve in water because they are hydrophobic.

Water, which not only dissolves many compounds but also dissolves more substances than any other liquid, is considered the universal solvent. A polar molecule with partially-positive and negative charges, it readily dissolves ions and polar molecules. Water is therefore referred to as a solvent: a substance capable of dissolving other polar molecules and ionic compounds. The charges associated with these molecules form hydrogen bonds with water, surrounding the particle with water molecules. This is referred to as a sphere of hydration, or a hydration shell, and serves to keep the particles separated or dispersed in the water.

When ionic compounds are added to water, individual ions interact with the polar regions of the water molecules during the dissociation process, disrupting their ionic bonds. Dissociation occurs when atoms or groups of atoms break off from molecules and form ions. Consider table salt (NaCl, or sodium chloride): when NaCl crystals are added to water, the molecules of NaCl dissociate into Na+ and Cl ions, and spheres of hydration form around the ions. The positively-charged sodium ion is surrounded by the partially-negative charge of the water molecule’s oxygen; the negatively-charged chloride ion is surrounded by the partially-positive charge of the hydrogen in the water molecule.

Dissociation of NaCl in water: When table salt (NaCl) is mixed in water, spheres of hydration form around the ions.

Since many biomolecules are either polar or charged, water readily dissolves these hydrophilic compounds. Water is a poor solvent, however, for hydrophobic molecules such as lipids. Nonpolar molecules experience hydrophobic interactions in water: the water changes its hydrogen bonding patterns around the hydrophobic molecules to produce a cage-like structure called a clathrate. This change in the hydrogen-bonding pattern of the water solvent causes the system’s overall entropy to greatly decrease, as the molecules become more ordered than in liquid water. Thermodynamically, such a large decrease in entropy is not spontaneous, and the hydrophobic molecule will not dissolve.