7.4 pH Scale and Neutralization Reactions

The pH scale is a crucial tool for measuring acidity and basicity in solutions. It ranges from 0 to 14, with 7 being neutral. Understanding pH helps us classify solutions as acidic, basic, or neutral, which is essential in various fields from chemistry to biology.

Neutralization reactions occur when acids and bases combine, forming water and a salt. This process is vital in many applications, including antacids neutralizing stomach acid. Titration, a technique used to determine unknown concentrations, relies on neutralization reactions and precise measurements.

pH and Solution Types

Understanding the pH Scale and Ion Concentrations

Top images from around the web for Understanding the pH Scale and Ion Concentrations

• 

  Relative Strengths of Acids and Bases | General Chemistry View original

  Is this image relevant?

• 

  pH and pOH | General Chemistry View original

  Is this image relevant?

• 

  2.4: The pH Scale - Biology LibreTexts View original

  Is this image relevant?

• 

  Relative Strengths of Acids and Bases | General Chemistry View original

  Is this image relevant?

• 

  pH and pOH | General Chemistry View original

  Is this image relevant?

1 of 3

Top images from around the web for Understanding the pH Scale and Ion Concentrations

• 

  Relative Strengths of Acids and Bases | General Chemistry View original

  Is this image relevant?

• 

  pH and pOH | General Chemistry View original

  Is this image relevant?

• 

  2.4: The pH Scale - Biology LibreTexts View original

  Is this image relevant?

• 

  Relative Strengths of Acids and Bases | General Chemistry View original

  Is this image relevant?

• 

  pH and pOH | General Chemistry View original

  Is this image relevant?

1 of 3

• pH scale measures acidity or basicity of a solution ranging from 0 to 14
• Hydrogen ion concentration [H+] determines acidity of a solution
• Hydroxide ion concentration [OH-] determines basicity of a solution
• pH calculated using the formula $pH = -log[H+]$
• pOH calculated using the formula $pOH = -log[OH-]$
• Relationship between pH and pOH expressed as $pH + pOH = 14$
• Higher [H+] results in lower pH, indicating higher acidity
• Higher [OH-] results in higher pH, indicating higher basicity

Classifying Solutions Based on pH

• Neutral solution has equal concentrations of H+ and OH- ions
  • pH equals 7 at room temperature (25°C)
  • Pure water serves as an example of a neutral solution
• Acidic solution contains more H+ ions than OH- ions
  • pH falls below 7
  • Lemon juice (pH ~2) and vinegar (pH ~3) represent common acidic solutions
• Basic solution contains more OH- ions than H+ ions
  • pH exceeds 7
  • Household ammonia (pH ~11) and baking soda solution (pH ~9) exemplify basic solutions
• pH indicators change color based on the solution's pH
  • Litmus paper turns red in acidic solutions and blue in basic solutions
  • Universal indicator displays a spectrum of colors for different pH values

Neutralization and Titration

Neutralization Reactions and Salt Formation

• Neutralization reaction occurs when an acid and a base combine
• Produces water and a salt as products
• General equation: $Acid + Base \rightarrow Salt + Water$
• Hydrochloric acid and sodium hydroxide react to form sodium chloride and water $HCl + NaOH \rightarrow NaCl + H_2O$
• Salt formed depends on the specific acid and base used in the reaction
• Neutralization reactions release heat, classified as exothermic
• Applications include antacids neutralizing stomach acid (calcium carbonate neutralizing hydrochloric acid)

Titration Process and Techniques

• Titration determines the concentration of an unknown acid or base solution
• Involves adding a known concentration of base (or acid) to a known volume of acid (or base)
• Uses a buret to precisely measure the volume of titrant added
• Endpoint indicated by a sudden change in pH or color of an indicator
• Equivalence point occurs when moles of acid equal moles of base
• Calculations use the formula $M_aV_a = M_bV_b$ where M represents molarity and V represents volume
• Phenolphthalein serves as a common indicator, turning pink in basic solutions
• Methyl orange provides another indicator option, changing from red in acidic to yellow in basic solutions

Buffer Solutions and Their Function

• Buffer solutions 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
• Henderson-Hasselbalch equation describes buffer behavior: $pH = pK_a + log(\frac{[A^-]}{[HA]})$
• Common buffer systems include acetic acid/acetate ion and ammonia/ammonium ion
• Biological systems rely on buffers to maintain optimal pH (blood pH maintained at ~7.4)
• Buffer capacity refers to the amount of acid or base a buffer can neutralize before significant pH change
• Preparing buffers involves combining appropriate concentrations of weak acid/base and its salt