Acid-Base Chemistry: pH, Titration, and Buffer Systems
Acids and bases touch nearly every aspect of chemistry and biology. Your stomach uses hydrochloric acid to digest food. Your blood relies on a precise buffer system to maintain pH between 7.35 and 7.45 — deviation of 0.1 units can be fatal. Industrial processes from fertilizer production to petroleum refining depend on acid-base chemistry.
The sour taste of lemons, the slippery feel of soap, the fizz of antacids in water — these everyday experiences have a unified chemical explanation. Understanding acids and bases gives you insight into processes ranging from digestion to industrial chemistry.
Defining Acids and Bases
The Arrhenius Definition
Svante Arrhenius defined acids as substances that produce hydrogen ions (H+) when dissolved in water and bases as substances that produce hydroxide ions (OH-). Hydrochloric acid (HCl) dissociates to H+ and Cl-. Sodium hydroxide (NaOH) dissociates to Na+ and OH-.
This definition works well for aqueous solutions but fails for bases that contain no hydroxide, like ammonia. Ammonia produces OH- in water but contains none itself: NH3 + H2O → NH4+ + OH-.
The Brønsted-Lowry Definition
The Brønsted-Lowry definition broadens the concept: acids are proton (H+) donors, and bases are proton acceptors. This explains ammonia’s basic behavior — it accepts a proton from water. Every Brønsted-Lowry acid has a conjugate base (what remains after donating a proton), and every base has a conjugate acid.
In the reaction HCl + H2O → Cl- + H3O+, HCl donates a proton to water. HCl is the acid, H2O is the base, Cl- is the conjugate base, and H3O+ is the conjugate acid. Water is amphoteric — it can act as either acid or base depending on the other substance.
The Lewis Definition
G. N. Lewis provided the most general definition: acids are electron-pair acceptors, and bases are electron-pair donors. This encompasses reactions that do not involve protons at all. Boron trifluoride (BF3) is a Lewis acid because it can accept an electron pair. Ammonia is a Lewis base because it donates its lone pair.
The pH Scale
The concentration of H+ in pure water at 25°C is 1.0 × 10^-7 M. The pH scale, defined as pH = -log[H+], makes these tiny numbers manageable. Pure water has pH 7.0. Acidic solutions have pH < 7. Basic solutions have pH > 7.
Each pH unit represents a tenfold change in [H+]. A solution of pH 3 is ten times more acidic than pH 4 and one hundred times more acidic than pH 5. This logarithmic scale means small pH changes represent large changes in acidity.
Strong acids like HCl and HNO3 dissociate completely in water. A 0.1 M HCl solution has [H+] = 0.1 M, so pH = 1.0. Weak acids like acetic acid (CH3COOH) dissociate only partially. A 0.1 M acetic acid solution has pH about 2.9, corresponding to only about 1% dissociation.
pOH and Kw
pOH = -log[OH-]. In aqueous solution at 25°C, pH + pOH = 14. The ion product of water, Kw = [H+][OH-] = 1.0 × 10^-14 at 25°C, governs the relationship between H+ and OH- concentrations. This relationship is crucial for solution chemistry calculations.
Strong and Weak Acids and Bases
Strong acids and bases dissociate completely. The six common strong acids are HCl, HBr, HI, HNO3, H2SO4, and HClO4. Strong bases are the group 1 and group 2 metal hydroxides: NaOH, KOH, Ca(OH)2, and Ba(OH)2.
Weak acids and bases dissociate only partially in water. Acetic acid (CH3COOH), carbonic acid (H2CO3), and phosphoric acid (H3PO4) are common weak acids. Ammonia (NH3) is a common weak base.
The acid dissociation constant Ka measures acid strength. Larger Ka means stronger acid. Acetic acid has Ka = 1.8 × 10^-5, while hydrofluoric acid has Ka = 6.8 × 10^-4. pKa = -log Ka, so smaller pKa means stronger acid.
Acid-Base Titration
Titration is the controlled addition of a base to an acid (or vice versa) to determine concentration. The equivalence point occurs when moles of H+ equal moles of OH-. An indicator changes color at the endpoint, which should be as close as possible to the equivalence point.
Titration Curves
Strong acid-strong base titrations show a sharp pH change near the equivalence point (pH 7). The curve is nearly flat until just before the equivalence point, then jumps dramatically over a few drops of titrant.
Weak acid-strong base titrations have the equivalence point at pH > 7 because the conjugate base of the weak acid hydrolyzes (reacts with water to produce OH-). The starting pH is higher than for a strong acid of the same concentration. A buffer region appears before the equivalence point where the solution resists pH change.
Weak base-strong acid titrations have the equivalence point at pH < 7. The half-equivalence point equals pKa for weak acid titrations and pKb for weak base titrations.
Buffer Solutions
Buffers resist pH change when small amounts of acid or base are added. They consist of a weak acid and its conjugate base (or a weak base and its conjugate acid). Buffer capacity depends on the amounts of the two components.
The Henderson-Hasselbalch equation relates buffer pH to component concentrations: pH = pKa + log([conjugate base]/[weak acid]). This equation allows precise buffer preparation.
Biological Buffers
Blood is buffered primarily by the carbonic acid/bicarbonate system: H2CO3/HCO3-. This system maintains blood pH at 7.4. The equilibrium H2O + CO2 ⇌ H2CO3 ⇌ H+ + HCO3- connects breathing rate to blood pH — hyperventilation lowers CO2, raising pH (respiratory alkalosis).
Phosphate buffers operate inside cells where pH changes would disrupt enzyme function. Proteins themselves act as buffers through their ionizable amino acid side chains.
Applications of Acid-Base Chemistry
Acid-base chemistry has countless practical applications. Antacids neutralize excess stomach acid. Agricultural lime (CaCO3) neutralizes acidic soil. Industrial wastewater pH must be adjusted before discharge. The ocean’s pH has dropped 0.1 units since the industrial revolution due to CO2 absorption — ocean acidification threatens marine organisms that build calcium carbonate shells.
Measuring pH is essential in food production, water treatment, pharmaceutical manufacturing, and swimming pool maintenance. Electrochemistry also involves acid-base chemistry in electrochemical cells and pH sensors.
Polyprotic Acids and Amphoteric Substances
Polyprotic acids can donate more than one proton. Sulfuric acid (H2SO4) is diprotic, dissociating in two steps. The first dissociation is complete (strong acid), but the second (HSO4- → H+ + SO4^2-) is partial with Ka2 = 0.012. Phosphoric acid (H3PO4) is triprotic, with three successive dissociations, each weaker than the previous.
Each dissociation step has its own Ka value. For phosphoric acid: Ka1 = 7.5 × 10^-3, Ka2 = 6.2 × 10^-8, Ka3 = 4.8 × 10^-13. The vast differences mean that at a given pH, the predominant species can be predicted. At pH 7, H2PO4- and HPO4^2- dominate, making phosphate an excellent buffer near neutral pH.
Zwitterions and Amino Acids
Amino acids contain both acidic (carboxyl) and basic (amino) functional groups. In neutral solution, the carboxyl loses a proton and the amino gains one, forming a zwitterion with both positive and negative charges. The isoelectric point is the pH at which the molecule has no net charge.
The zwitterionic nature of amino acids makes them soluble in water and gives them buffer capacity. Proteins, as polymers of amino acids, inherit these acid-base properties. The buffering capacity of proteins in blood and cells is essential for maintaining physiological pH.
Industrial Acid-Base Processes
The Contact process for sulfuric acid production is the largest-volume chemical manufacturing process in the world. Sulfuric acid is a diprotic acid used in fertilizer production, petroleum refining, mineral processing, and chemical synthesis. Global production exceeds 250 million tons annually, making it the most produced chemical by volume.
The Solvay process for sodium carbonate uses acid-base chemistry at every stage. Carbon dioxide is absorbed into ammoniated brine to precipitate sodium bicarbonate, which is then calcined to sodium carbonate. The process recycles ammonia and CO2, demonstrating elegant industrial chemistry.
Frequently Asked Questions
What pH is considered neutral? pH 7.0 at 25°C. Neutral pH changes with temperature because Kw varies — pH 7.0 is neutral only at 25°C.
How do buffers resist pH change? Buffers contain both a weak acid and its conjugate base. Added acid reacts with the base component. Added base reacts with the acid component. This consumes the added H+ or OH- with minimal pH change.
What is the difference between a strong acid and a concentrated acid? Strength refers to dissociation extent (complete vs. partial). Concentration refers to amount of acid per volume. A concentrated weak acid (like concentrated acetic acid) can be more acidic than a dilute strong acid.
Why do weak acids have higher pH than strong acids of the same concentration? Weak acids dissociate only partially, producing fewer H+ ions than strong acids at the same concentration.
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