All Study Guides Biological Chemistry I Unit 2
🔬 Biological Chemistry I Unit 2 – Water, pH, and BuffersWater, pH, and buffers form the foundation of biological chemistry. These concepts explain how life processes function at the molecular level. Understanding water's unique properties, pH scales, and buffer systems is crucial for grasping how cells maintain homeostasis and carry out essential biochemical reactions.
This unit covers water's polarity and hydrogen bonding, pH measurement, acids and bases, and buffer systems. It explores how these factors influence biological processes, from enzyme function to cellular compartmentalization. The knowledge gained here is essential for comprehending more complex biochemical interactions in living organisms.
Got a Unit Test this week? we crunched the numbers and here's the most likely topics on your next test Key Concepts
Water's unique properties arise from its polar nature and hydrogen bonding capabilities
Polarity results from uneven distribution of electrons between hydrogen and oxygen atoms
Hydrogen bonding occurs between partially positive hydrogen and partially negative oxygen of neighboring water molecules
pH is a logarithmic scale that measures the concentration of hydrogen ions (H+) in a solution
Calculated as the negative log of the hydrogen ion concentration: p H = − l o g [ H + ] pH = -log[H+] p H = − l o g [ H + ]
pH values range from 0 to 14, with 7 being neutral, below 7 acidic, and above 7 basic
Acids are proton donors that increase the concentration of H+ ions in a solution
Strong acids (hydrochloric acid) dissociate completely in water
Weak acids (acetic acid) only partially dissociate in water
Bases are proton acceptors that decrease the concentration of H+ ions in a solution
Strong bases (sodium hydroxide) dissociate completely in water
Weak bases (ammonia) only partially dissociate in water
Buffer systems resist changes in pH when small amounts of acid or base are added
Consist of a weak acid and its conjugate base, or a weak base and its conjugate acid
Important in maintaining stable pH in biological systems (blood, cytoplasm)
Water's Unique Properties
High specific heat capacity enables water to absorb and release heat slowly
Helps regulate temperature in living organisms and Earth's climate
Strong cohesive and adhesive forces due to hydrogen bonding
Cohesion allows water to maintain column strength in xylem vessels of plants
Adhesion enables capillary action, which is crucial for water transport in plants and soil
High surface tension resulting from cohesive forces among water molecules
Allows some insects (water striders) to walk on water's surface
Universal solvent capabilities due to polarity and ability to form hydrogen bonds
Dissolves a wide range of polar and ionic compounds (salts, sugars, proteins)
Essential for transport of nutrients and waste products in living organisms
Density anomaly: water is most dense at 4°C and less dense as a solid (ice)
Ice floats on liquid water, insulating aquatic life in cold climates
pH Basics
pH is a measure of the acidity or basicity of a solution
Acidic solutions have a pH below 7 and a higher concentration of H+ ions
Basic solutions have a pH above 7 and a lower concentration of H+ ions
Neutral solutions have a pH of 7 and equal concentrations of H+ and OH- ions
The pH scale is logarithmic, meaning each unit change represents a tenfold difference in H+ concentration
A solution with pH 5 has 10 times more H+ ions than a solution with pH 6
A solution with pH 9 has 100 times fewer H+ ions than a solution with pH 7
The concentration of H+ and OH- ions in a solution are inversely related
In pure water at 25°C, [ H + ] [ O H − ] = 1.0 × 1 0 − 14 [H+][OH-] = 1.0 \times 10^{-14} [ H + ] [ O H − ] = 1.0 × 1 0 − 14 (the ion product constant of water, Kw)
As [H+] increases, [OH-] decreases proportionally to maintain Kw
Acids and Bases
Arrhenius definition: acids produce H+ ions in solution, while bases produce OH- ions
Limited to aqueous solutions and does not explain all acid-base behaviors
Brønsted-Lowry definition: acids are proton donors, while bases are proton acceptors
Accounts for acid-base reactions in both aqueous and non-aqueous solutions
Conjugate acid-base pairs differ by a single proton (H+)
Lewis definition: acids are electron pair acceptors, while bases are electron pair donors
Broadest definition, encompassing both Arrhenius and Brønsted-Lowry concepts
Explains the basicity of compounds without OH- ions (ammonia)
Acid and base strength depend on the extent of dissociation in solution
Strong acids and bases dissociate completely (hydrochloric acid, sodium hydroxide)
Weak acids and bases only partially dissociate (acetic acid, ammonia)
The acid dissociation constant (Ka) quantifies the strength of an acid
Higher Ka values indicate stronger acids and more complete dissociation
Buffer Systems
Buffers minimize pH changes when small amounts of acid or base are added to a solution
Consist of a weak acid and its conjugate base, or a weak base and its conjugate acid
The weak acid or base component neutralizes added base or acid, respectively
The Henderson-Hasselbalch equation relates pH, pKa, and the ratio of base to acid concentrations in a buffer system:
p H = p K a + l o g [ b a s e ] [ a c i d ] pH = pK_a + log\frac{[base]}{[acid]} p H = p K a + l o g [ a c i d ] [ ba se ]
Buffer capacity is the amount of acid or base that can be added before the pH changes significantly
Determined by the concentrations of the buffer components and the pKa of the acid or base
Buffers are most effective when the pH is within ±1 unit of the pKa
Preparing a buffer with a specific pH involves selecting an acid-base pair with a pKa close to the desired pH
The ratio of base to acid concentrations is adjusted according to the Henderson-Hasselbalch equation
Common biological buffers include bicarbonate (HCO3-/CO2) and phosphate (H2PO4-/HPO4-2) systems
Bicarbonate buffer helps maintain blood pH around 7.4
Phosphate buffer helps maintain intracellular pH around 7.2
Biological Importance
Maintaining stable pH is crucial for proper functioning of biological systems
Enzymes have optimal pH ranges for activity; deviations can lead to loss of function
Protein structure and solubility are sensitive to pH changes
Membrane potential and ion transport are affected by pH gradients
Blood pH is tightly regulated between 7.35 and 7.45 by buffer systems, respiratory control, and renal control
Bicarbonate buffer is the primary blood buffer, with hemoglobin and plasma proteins also contributing
Respiratory control adjusts blood pH by regulating CO2 levels through breathing rate
Renal control adjusts blood pH by excreting excess H+ or HCO3- ions in urine
Cellular compartments have specific pH ranges optimized for their functions
Cytoplasm: pH 7.2, favorable for most enzymatic reactions
Lysosomes: pH 4.5-5.0, optimal for acid hydrolase enzymes involved in digestion
Mitochondrial matrix: pH 7.8, supports efficient ATP synthesis during oxidative phosphorylation
pH gradients across membranes are harnessed for energy production and transport processes
Proton motive force drives ATP synthesis in mitochondria and chloroplasts
pH gradients facilitate secondary active transport of nutrients (glucose, amino acids) into cells
Lab Techniques and Measurements
pH meters measure the electrical potential difference between a reference electrode and a glass electrode in a solution
The potential difference is proportional to the pH of the solution
Calibration with standard buffer solutions (pH 4, 7, 10) is necessary for accurate measurements
pH paper and indicators change color depending on the pH of the solution
Litmus paper turns red in acidic solutions and blue in basic solutions
Universal indicator displays a range of colors corresponding to pH values from 0 to 14
Specific indicators (phenolphthalein, methyl orange) have distinct color changes at certain pH ranges
Titration is used to determine the concentration of an acid or base in a solution
A known concentration of base or acid (titrant) is gradually added to the unknown solution (analyte)
The endpoint is reached when the pH of the analyte changes abruptly, often indicated by a color change
The concentration of the analyte is calculated based on the volume and concentration of the titrant added
Spectrophotometric methods can measure pH-dependent changes in absorbance
Some compounds (phenol red) exhibit different colors and absorbance spectra at different pH values
The ratio of absorbance at two wavelengths can be used to calculate the pH of a solution
Real-World Applications
Environmental monitoring: pH measurements assess the health of aquatic ecosystems and soil quality
Acidification of lakes and streams due to acid rain can harm aquatic life
Soil pH affects nutrient availability and plant growth; lime is used to raise pH in acidic soils
Food and beverage industry: pH control ensures product safety, quality, and shelf life
Low pH in fermented foods (yogurt, pickles) prevents the growth of spoilage bacteria
Adjusting pH during cheese making influences texture and flavor development
Pharmaceutical industry: pH considerations are crucial in drug formulation and delivery
Drug solubility and absorption are pH-dependent; oral medications often have pH-sensitive coatings
Buffer systems maintain the stability and efficacy of liquid medications
Wastewater treatment: pH adjustment is necessary for effective removal of contaminants
Alkaline conditions (pH 10-11) promote the precipitation of heavy metals
Neutral pH is maintained in biological treatment stages to support microbial activity
Medical diagnostics: pH abnormalities in body fluids can indicate underlying health conditions
Blood pH outside the normal range (7.35-7.45) can signify respiratory or metabolic disorders
Urine pH can provide insights into kidney function, urinary tract infections, and metabolic status