pH: Ácidos y Bases
Understanding Acids and Bases
Introduction to Acids
- The class will explore the concept of pH, starting with basic definitions of acids and bases. The term "acid" originates from the Latin word "acidus," meaning sour, which reflects the characteristic sour taste of acidic substances.
- When an acid reacts with litmus paper, it turns red. Additionally, acids react with metals to release hydrogen ions (protons), resulting in a tactile sensation that feels watery.
Characteristics of Bases
- Bases, also known as alkalis, derive their name from the Greek word "alkali," meaning ash. They are characterized by a bitter taste and turn red litmus paper blue when tested.
- Touching a basic substance like bleach gives a slippery or soapy feeling, indicating its alkaline nature.
Definitions by Lowry and Bronsted
- In 1923, English chemist Lowry and Danish chemist Bronsted independently defined acids as substances that donate protons (H⁺ ions), while bases accept these protons.
- This definition has biochemical applications; for an acid to exist in a medium, there must be a corresponding base present to accept its protons.
Water's Dual Role
- Water exhibits properties of both an acid and a base: it can donate protons or accept them. For example:
- When reacting with hydrochloric acid (HCl), water acts as a base by accepting protons from HCl.
- Conversely, when interacting with ammonia (NH₃), water behaves as an acid by donating protons.
Reactions Involving Water
- A reaction between two water molecules demonstrates that one can act as an acid while the other acts as a base simultaneously—one donates protons while the other accepts them.
- These reactions illustrate how acids and bases behave in aqueous solutions without explicitly including water molecules in biochemical equations for clarity.
Auto-dissociation of Water
- Water can dissociate into hydroxide ions (OH⁻) and hydronium ions (H₃O⁺). This self-ionization is termed auto-dissociation.
- At equilibrium during this process, we define the constant K_w , representing the product of concentrations of hydronium and hydroxide ions ( K_w = [H^+][OH^-] ), which equals 1 times 10^-14 .
Implications on Solution Properties
Understanding pH and Hydroxide Concentration
The Relationship Between Protons and Hydroxide Ions
- A neutral solution occurs when the number of protons equals the number of hydroxide ions, resulting in equal concentrations.
- A basic solution is present when there are more hydroxide ions than protons, indicating a higher concentration of hydroxide compared to protons.
- The natural concentration of these substances is typically very low, making it challenging to express values accurately using decimals or exponential notation.
Introduction of pH by Sorensen
- In 1909, Danish chemist Sorensen proposed a logarithmic transformation for expressing proton concentration, termed "pH," defined as the negative logarithm (base 10) of proton concentration.
- This transformation converts fractional numbers into positive integers; thus, a higher concentration of hydrogen ions results in a lower pH value.
Calculating pH in Pure Water
- The dissociation constant of water is expressed as the product of hydrogen ion and hydroxide ion concentrations, valued at 1 times 10^-14.
- In pure water, for every hydrogen ion produced, an equivalent amount of hydroxide ion exists; therefore, their concentrations are equal.
Understanding pH Values
- The calculated value for pure water's pH is 7 at 25 degrees Celsius. This indicates that both hydrogen and hydroxide concentrations are equal at this point.
- When calculating either pH or its complementary value (pOH), they will always sum to 14 in aqueous solutions.
Observations from the pH Table
- The table illustrates that as proton concentration increases (moving upwards), the corresponding pH decreases; hence lower values indicate higher acidity.