buffer state

Examples of buffer state in the following topics:

• pH, Buffers, Acids, and Bases

  • By convention, scientists refer to hydrogen ions and their concentration as if they were free in this state in liquid water.
  • Buffers are the key.
  • Without this buffer system, the body's pH would fluctuate enough to jeopardize survival.
  • Antacids, which combat excess stomach acid, are another example of buffers.
  • Explain the composition of buffer solutions and how they maintain a steady pH

• Buffer Range and Capacity

  • A buffer's capacity is the pH range where it works as an effective buffer, preventing large changes in pH upon addition of an acid or base.
  • A titration curve visually demonstrates buffer capacity.
  • This is the buffer zone.
  • However, once the curve extends out of the buffer region, it will increase tremendously when a small amount of acid or base added to the buffer system.
  • Discuss correlation between the pKa of the conjugate acid of a buffer solution and the effective range of the corresponding buffer.

• The Common Ion Effect

  • Chemical equilibrium is the chemical state where there are no net physical or chemical changes between the reactant and the products of a reaction.
  • Solubility equilibrium refers to the state of chemical equilibrium between a chemical compound in the solid state and a solution composed of that dissolved compound.
  • A buffer solution is composed of a weak acid and its conjugate base, or a weak base and its conjugate acid.
  • Addition of excess ions will alter the pH of the buffer solution.
  • In the case of an an acidic buffer, the hydrogen ion concentration decreases, and the resulting solution is less acidic than a solution containing the pure weak acid.

• Chemical Buffer Systems

  • Chemical buffers such as bicarbonate and ammonia help keep blood pH in the narrow range compatible with life.
  • One example of a buffer solution found in nature is blood.
  • Several buffering agents that reversibly bind hydrogen ions and impede any change in pH exist.
  • Extracellular buffers include bicarbonate and ammonia, whereas proteins and phosphate act as intracellular buffers.
  • Distinguish between buffer solutions, ventilation, and renal function as buffer systems to control acid-base balance

• Calculating Changes in a Buffer Solution

  • The changed pH of a buffer solution in response to the addition of an acid or a base can be calculated.
  • These solutions are known as buffers.
  • What would be the pH of the sodium hydroxide solution without the buffer?
  • Solving for the buffer pH after 0.0020 M NaOH has been added:
  • It also shows the importance of using high buffer component concentrations so that the buffering capacity of the solution is not exceeded.

• Preparing a Buffer Solution with a Specific pH

  • For example, blood in the human body is a buffer solution.
  • Buffer solutions are necessary in a wide range of applications.
  • There are a couple of ways to prepare a buffer solution of a specific pH.
  • Both solutions must contain the same buffer concentration as the concentration of the buffer in the final solution.
  • To get the final buffer, add one solution to the other while monitoring the pH.

• Relative Amounts of Acid and Base

  • A buffer's pH changes very little when a small amount of strong acid or base is added to it.
  • Suppose you wish to prepare a buffer solution to keep the pH at 4.30.
  • What amount of acid and base should you use to create the buffer?
  • This is due to the change that occurs when another acid or base is added to the buffer.
  • The more the ratio needs to differ to achieve the desired pH, the less effective the buffer.

• Absolute Concentrations of the Acid and Conjugate Base

  • A buffer's pH changes very little when a small amount of strong acid or base is added to it.
  • A concentrated buffer can neutralize more added acid or base than a dilute buffer, because it contains more acid/conjugate base.
  • However, any buffer will lose its effectiveness if too much strong acid or base is added.
  • Therefore, the pH for the buffer with an acid/base concentration of 0.7/0.6M is 4.68.
  • Therefore, the pH of the weaker buffer before the addition of HCl is the same.

• Acidosis

  • The term "acidemia" describes the state of low blood pH, while acidosis is used to describe the processes leading to these states.
  • Metabolic acidosis is compensated for in the lungs, as increased exhalation of carbon dioxide promptly shifts the buffering equation to reduce metabolic acid.
  • The Henderson-Hasselbalch equation is useful for calculating blood pH, because blood is a buffer solution.
  • The amount of metabolic acid accumulating can also be quantitated by using buffer base deviation, a derivative estimate of the metabolic as opposed to the respiratory component.
  • Compensation occurs if respiratory acidosis is present, and a chronic phase is entered with partial buffering of the acidosis through renal bicarbonate retention.

• The Acid Dissociation Constant

  • The equilibrium constant K = 1 states that there will be 50 percent products and 50 percent reactants.
  • An understanding of Ka is also essential for working with buffers; the design of these solutions depends on a knowledge of the pKa values of their components.
  • Buffers are used whenever there is a need to fix the pH of a solution at a particular value.
  • Buffering is an essential part of in-vitro biochemical studies and acid-base physiology and plays a key role in analytical chemistry.
  • Compared with an aqueous solution, the pH of a buffer solution is relatively insensitive to the addition of a small amount of strong acid or strong base.