Chapter 1. Water, Aqueous Solutions, pH, and Buffers

Water, Aqueous Solutions, pH, and Buffers

Water is an amazing molecule. Life began in water and evolved there for three billion years before spreading on land. All modern life on Earth depends on water. Most of your body (~70%) is made of water. Since the properties of water are closely related to its structure, we start this lab with a discussion of water’s molecular structure.

Water, H₂O, is made up of two hydrogen atoms and one oxygen atom connected by a polar covalent bond. A polar covalent bond is a type of chemical bond. Chemical bonds are attractions that hold atoms together. Covalent bonds are formed by sharing electrons. A single bond is the sharing of one pair of electrons. In the case of a water molecule, oxygen has an especially high attraction to electrons—it has a high electronegativity. This means that it pulls the shared electrons closer to it than to hydrogen, giving oxygen a partial negative charge (δ–) while the two hydrogens have partial positive charges (δ+). A covalent bond where the electrons are not evenly shared is called a polar covalent bond.

Opposite charges attract. Therefore, the partial negative charge on the oxygen of one water molecule is attracted to the partial positive charge on a hydrogen atom of another water molecule. This attraction causes the formation of temporary, weak hydrogen bonds between water molecules in liquid water. A hydrogen bond is much weaker than a covalent bond. Each water molecule can form four intermolecular hydrogen bonds.

Water is a substance with many unique properties. Much of what water can do that other liquids cannot is based upon water’s polarity. Water is a polar molecule (remember, partially negative O and partially positive H). Properties of water important for this lab are (1) water is a liquid at standard temperature and pressure and (2) water is a good solvent.

A solution is a liquid homogeneous mixture of two or more substances. A solvent is the dissolving agent of a solution. A solute is a substance that is dissolved. An aqueous solution is a solution where water is the solvent. Biological chemistry is a chemistry of aqueous solutions. Here we will discuss two characteristics of aqueous solutions: concentrations and pH.

Solute Concentration and Molarity

Solute concentration is the measure of the amount of a solute in a solution. It can be expressed as percent by mass (20% solution of NaCl), which is a mass concentration (20 g/L of NaCl). We will use molar concentration (1 M/L). A Mole is an exact number of objects: 6.02 × 10²³ (Avogadro’s number). Molarity is the number of moles of solute per liter of solution. Atomic mass is the mass of an atom of a chemical element that is expressed in amu (atomic mass units). Molar mass is the mass of all the atoms of a molecule in gram per mole. The molar mass of an atom is equal to the atomic mass.

Acids, Bases, and pH Scale

Sometimes a water molecule dissociates: a hydrogen ion (H⁺) leaves one water molecule, creating a hydroxide ion (OH^-).

The hydrogen ion typically associates with another water molecule to form a hydronium ion (H₃O⁺) since the proton’s positive charge is attracted to the partial negative charge on H₂O’s oxygen atom. It happens rarely: in pure water, about 1 water molecule in 555 million is dissociated at any given time, producing a hydrogen ion/hydronium ion concentration of 10^-7 M.

[H⁺] = [OH^-] = 10^-7 M

Brackets are usually used to denote concentration. The [H⁺] in a solution are of critical interest to biologists and chemists because it has a major effect on chemical reactions and on H⁺ association with and dissociation from other molecules. A wide range of hydrogen ion concentrations is possible, so scientists use the logarithmic pH scale to describe [H⁺]. pH is a negative logarithm (base 10) of the hydrogen ion concentration:

pH = – log [H⁺].

The pH of pure water equals 7; this is called the neutral pH. For all aqueous solutions

pH + pOH = 14 (pOH = – log [OH^-]).

This means that pH = pOH only at pH 7. If you add H⁺ to a solution (decreasing the pH), the added protons associate with some of the OH^-, forming water and causing a decrease in the [OH^-] (which increases the pOH). Note that the pH scale is logarithmic, so a one unit change in pH equals a 10-fold change in [H⁺]. Since pH equals the negative logarithm of [H⁺], a one unit increase in pH equals a 10-fold decrease in [H⁺].

Many substances besides water also dissociate H⁺ in solution; these proton donors are called acids and they increase the [H⁺] in a solution, causing the solution to have a pH < 7. Some common acids are lemon juice, vinegar, and cola. Substances that bind H⁺ or release OH^- (which in turn binds H⁺ to form H₂O) are called bases. Bases decrease the [H⁺] of a solution, causing the solution to have a pH > 7. Some common bases are ammonia and baking soda (bicarbonate).

Buffers can both donate protons and accept protons (or generate hydroxide ions) - they resist changes in pH. A buffer tends to keep a solution at a constant pH by donating protons when protons are removed from the solution and accepting protons (or releasing OH^-) when protons are added to the solution. However, if enough protons are added or removed from the solution, the buffering capacity will eventually be exhausted and the pH will change. Bicarbonate ion, or hydrogencarbonate, (HCO₃^-) is an important buffer that helps maintain your blood pH (normal blood pH is 7.4).

Buffer Action of Bicarbonate Solution

We will be using bicarbonate ions from baking soda (NaHCO₃) to investigate buffer action. You will compare pH changes in pure water and in bicarbonate solution (buffer) following the gradual addition of lemon juice (an acid).