Difficulty: Introductory | Prerequisites: Basic algebra, logarithms, scientific notation.
This topic sits at the heart of general chemistry: understanding what makes a solution acidic or basic, how to quantify that on the pH scale, and how to calculate equilibrium constants for weak acids and bases. It connects directly to buffer systems, titrations, and solubility, all of which build on the ideas here. You should already be comfortable with molarity, the concept of chemical equilibrium, and basic logarithm operations. If the word "equilibrium" still feels fuzzy, revisit your notes on Le Chatelier's principle before diving in.
pH measures how acidic or basic a solution is, on a scale from 0 to 14 (7 is neutral). You calculate it from the hydronium ion concentration using pH = -log[H3O+]. Strong acids and bases dissociate completely, so their pH comes straight from their molarity; weak acids and bases only partially dissociate, so you need the equilibrium constant (Ka or Kb) and an ICE table to find the pH.
pH
The negative base-10 logarithm of the hydronium ion concentration: pH = -log[H3O+]. Values below 7 are acidic, above 7 are basic, and 7 is neutral. Think of it as a compact way to express very small ion concentrations on a human-friendly scale.
pOH
The negative base-10 logarithm of the hydroxide ion concentration: pOH = -log[OH-]. Related to pH by the expression pH + pOH = 14 (at 25 degrees C). In simple terms, it does for bases what pH does for acids.
Acid
A substance that increases the concentration of H+ (or H3O+) ions in aqueous solution. Acids have a pH below 7.
Base
A substance that increases the concentration of OH- ions in aqueous solution. Bases have a pH above 7.
Arrhenius acid
A substance that produces hydrogen ions (H+) when dissolved in water. Example: HCl dissolves in water to release H+ and Cl-. Think of it as the older, water-only definition of an acid.
Arrhenius base
A substance that produces hydroxide ions (OH-) when dissolved in water. Example: NaOH dissolves to give Na+ and OH-.
Bronsted-Lowry acid
A proton (H+) donor in any reaction, not limited to water. Example: HCl donates a proton to water, forming H3O+. In simple terms, this is the broader, more modern definition that works in any solvent.
Bronsted-Lowry base
A proton (H+) acceptor. Example: NH3 accepts a proton from water to form NH4+.
Hydronium ion (H3O+)
The form hydrogen ions take in aqueous solution. When we write [H+], we really mean [H3O+].
Hydroxide ion (OH-)
The ion produced by bases in water. Its concentration determines pOH.
Strong acid
An acid that dissociates completely in water. All of the acid molecules break apart into ions. Examples: HCl, HNO3, H2SO4. Because dissociation is complete, [H3O+] equals the initial acid concentration.
Weak acid
An acid that only partially dissociates in water. An equilibrium exists between the undissociated acid (HA) and its ions (H+ and A-). Examples: HClO, H2CO3, CH3COOH.
Strong base
A base that dissociates completely in water. Examples: NaOH, KOH, Ba(OH)2.
Weak base
A base that only partially accepts protons in water, establishing an equilibrium. Example: NH3.
Ka (acid dissociation constant)
The equilibrium constant for the dissociation of a weak acid: Ka = [H+][A-] / [HA]. A larger Ka means a stronger (more dissociated) weak acid.
Kb (base dissociation constant)
The equilibrium constant for the reaction of a weak base with water: Kb = [BH+][OH-] / [B]. Related to Ka by the expression Ka x Kb = Kw.
Kw (ion product of water)
The product of [H3O+] and [OH-] in pure water at 25 degrees C: Kw = 1.0 x 10^-14. This constant links pH and pOH.
Conjugate acid-base pair
Two species that differ by exactly one proton. When an acid donates a proton, what remains is its conjugate base, and vice versa. Example: HClO / ClO-.
The pH scale runs from 0 to 14 at 25 degrees C.
pH < 7: acidic
pH = 7: neutral (pure water)
pH > 7: basic
pH is calculated from the hydronium ion concentration: pH = -log[H3O+].
pOH is calculated from the hydroxide ion concentration: pOH = -log[OH-].
At 25 degrees C, pH + pOH = 14 always holds. This means if you know one, you can find the other.
Strong acids dissociate completely, so [H3O+] equals the initial molarity of the acid.
Nitric acid (HNO3), 0.00125 M: pH = -log(0.00125) = 2.90
Hydrochloric acid (HCl), 4.55 M: pH = -log(4.55) = -0.66
Yes, pH can be negative when [H3O+] exceeds 1 M.
Strong bases dissociate completely, so [OH-] equals the initial molarity of the base (or twice that for diprotic bases like Ba(OH)2).
Potassium hydroxide (KOH), 0.00133 M: pOH = -log(0.00133) = 2.88, so pH = 14 - 2.88 = 11.1
Barium hydroxide (Ba(OH)2), 0.10 M: produces 0.20 M OH-, so pOH = -log(0.20) = 0.70, pH = 14 - 0.70 = 13.30
Arrhenius acid: produces H+ ions in water. Example: HCl -> H+ + Cl-.
Arrhenius base: produces OH- ions in water. Example: NaOH -> Na+ + OH-.
Limitation: only works in aqueous solution.
Bronsted-Lowry acid: donates a proton (H+) to another species. Example: HCl donates H+ to H2O, forming H3O+.
Bronsted-Lowry base: accepts a proton. Example: NH3 accepts H+ from H2O, forming NH4+.
This is the broader definition. Every Arrhenius acid is a Bronsted-Lowry acid, but not every Bronsted-Lowry base is an Arrhenius base (NH3 does not contain OH-, yet it is a base because it accepts protons).
Once you calculate the pH, you can classify a substance more precisely than just "acid" or "base."
Classification | pH Range | Dissociation | Example from Lab |
|---|---|---|---|
Strong acid | Well below 7, often < 3 | Complete | HCl (pH -0.66), HNO3 (pH 2.90) |
Weak acid | Below 7, typically 3 to 6 | Partial | HClO (pH 5.17), H2CO3 (pH 4.6) |
Acidic salt | Slightly below 7 | Salt of weak base + strong acid | NH4Cl (pH 6.05) |
Neutral | 7 | N/A | Pure water |
Basic salt | Slightly above 7 | Salt of strong base + weak acid | NaClO (pH 9.28), NaHCO3 (pH 8.31) |
Weak base | Above 7, typically 8 to 11 | Partial | NH3 (pH 10.2) |
Strong base | Well above 7, often > 11 | Complete | KOH (pH 11.1), Ba(OH)2 (pH 13.30) |
The key insight: salts are not neutral by default. A salt formed from a strong base and a weak acid (like sodium hypochlorite, NaClO) produces a basic solution because the conjugate base (ClO-) hydrolyses in water. A salt from a weak base and a strong acid (like ammonium chloride, NH4Cl) produces an acidic solution because the conjugate acid (NH4+) donates protons to water.
\text{pH} = -\log[\text{H}_3\text{O}^+]\text{pOH} = -\log[\text{OH}^-]\text{pH} + \text{pOH} = 14 \quad (\text{at } 25°\text{C})K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}K_b = \frac{[\text{BH}^+][\text{OH}^-]}{[\text{B}]}K_w = [\text{H}_3\text{O}^+][\text{OH}^-] = 1.0 \times 10^{-14}K_a \times K_b = K_wFor amphoteric species like the hydrogen carbonate ion (HCO3-):
\text{pH} = \frac{pK_{a1} + pK_{a2}}{2}Lactic acid (C3H6O3, Ka approx. 1.4 x 10^-4) is what makes sour milk taste sour. Lactobacilli ferment lactose into lactic acid, lowering the pH. The same acid builds up in your muscles during intense exercise.
Swimming pool chemistry relies on hypochlorous acid (HClO) and its conjugate base hypochlorite (ClO-). The pH of the pool determines how much of the chlorine is in the active, disinfecting HClO form versus the less effective ClO- form. Pool operators target pH 7.2 to 7.6 for this reason.
Blood pH is maintained between 7.35 and 7.45 by the carbonic acid / bicarbonate buffer system (H2CO3 / HCO3-), the same pair that appears in this lab. A shift of even 0.1 pH units outside that range can be life-threatening.
Students often assume all salts are neutral. They are not. A salt's pH depends on the relative strengths of the acid and base that formed it. NH4Cl is acidic (pH 6.05); NaClO is basic (pH 9.28).
Students sometimes think pH cannot be negative. It can. Any strong acid with a concentration above 1 M gives a negative pH (HCl at 4.55 M has pH = -0.66).
Confusing Ka with Kb: Ka describes acid dissociation, Kb describes base reaction with water. They are linked by Ka x Kb = Kw, but they are not interchangeable. When calculating the pH of a salt like NH4Cl, you use Ka for the NH4+ ion, not Kb for NH3.
Writing ICE tables for strong acids or strong bases. Strong acids and bases dissociate completely, so there is no equilibrium to set up. ICE tables are only for weak acids, weak bases, and salt hydrolysis reactions.
Expect to be asked to calculate the pH of a strong acid, a strong base, a weak acid, and a weak base given their concentrations and equilibrium constants.
Know when to use an ICE table (weak species) and when not to (strong species).
Be ready to classify a salt solution as acidic or basic based on the parent acid and base that formed it.
The relationship Ka x Kb = Kw is commonly tested, especially when you are given Ka for a weak acid and asked to find the pH of its conjugate base.
Phosphoric acid is polyprotic. The lab uses only the first dissociation (Ka1 = 7.5 x 10^-3), but exam questions may ask about the second or third dissociation.
For amphoteric species (like HCO3-), know the shortcut formula: pH = (pKa1 + pKa2) / 2.
True or false: A solution with pH 3 is more acidic than one with pH 5. (True. Lower pH = higher [H3O+] = more acidic.)
Fill in the blank: pH + pOH = _____ at 25 degrees C. (14)
True or false: You need an ICE table to calculate the pH of 0.1 M HCl. (False. HCl is a strong acid; pH = -log(0.1) = 1.)
Fill in the blank: Ka x Kb = _____. (Kw = 1.0 x 10^-14)
True or false: A solution of sodium acetate (the salt of a strong base and a weak acid) is acidic. (False. It is basic, because the acetate ion acts as a weak base in water.)
Q: A 0.045 M solution of lactic acid (C3H6O3) has a pH of 2.65. Calculate Ka.
A: [H3O+] = 10^-2.65 = 0.0022 M. The equilibrium concentration of undissociated acid is 0.045 - 0.0022 = 0.0428 M. Ka = (0.0022)^2 / 0.0428 = 1.13 x 10^-4.
Q: The Ka of hypochlorous acid (HClO) is 3.5 x 10^-8. Calculate the pH of a 0.00128 M solution.
A: Set up an ICE table. Ka = x^2 / 0.00128 (assuming x is small relative to 0.00128). x = sqrt(3.5 x 10^-8 x 0.00128) = 6.7 x 10^-6 M. pH = -log(6.7 x 10^-6) = 5.17.
Q: Calculate the pH of a 0.00125 M solution of sodium hypochlorite (NaClO). The Ka of HClO is 3.5 x 10^-8.
A: ClO- is the conjugate base of HClO. Kb = Kw / Ka = (1.0 x 10^-14) / (3.5 x 10^-8) = 2.9 x 10^-7. From the ICE table: x = sqrt(2.9 x 10^-7 x 0.00125) = 1.9 x 10^-5 M. pOH = -log(1.9 x 10^-5) = 4.72. pH = 14 - 4.72 = 9.28.
Q: Classify the following as acidic, basic, or neutral in solution: NH4Cl, NaClO, KOH.
A: NH4Cl is acidic (NH4+ is the conjugate acid of the weak base NH3). NaClO is basic (ClO- is the conjugate base of the weak acid HClO). KOH is strongly basic (complete dissociation of a strong base).
Q: Why does barium hydroxide at 0.10 M give a pOH of 0.70, not 1.00?
A: Ba(OH)2 is diprotic: each formula unit releases two OH- ions. So [OH-] = 2 x 0.10 = 0.20 M, and pOH = -log(0.20) = 0.70.
This material feeds directly into buffer solutions (covered next in most general chemistry sequences). A buffer is just a weak acid paired with its conjugate base, and the Henderson-Hasselbalch equation (pH = pKa + log([A-]/[HA])) is a rearrangement of the Ka expression you have been using here.
Titration curves also build on this: the equivalence point of a weak acid/strong base titration is not at pH 7, precisely because the conjugate base left in solution hydrolyses, the same process you see in the sodium hypochlorite and sodium cyanide calculations above.
Solubility equilibria (Ksp) use the same ICE-table logic, just applied to slightly soluble salts rather than acids and bases.
pH scale, pOH, hydronium ion, hydroxide ion, acid dissociation constant, Ka, Kb, Kw, ion product of water, Arrhenius acid, Arrhenius base, Bronsted-Lowry acid, Bronsted-Lowry base, proton donor, proton acceptor, strong acid, weak acid, strong base, weak base, conjugate acid, conjugate base, salt hydrolysis, acidic salt, basic salt, amphoteric, ICE table, equilibrium expression, lactic acid, hypochlorous acid, sodium hypochlorite, ammonium chloride, carbonic acid, phosphoric acid, barium hydroxide, sodium cyanide, Henderson-Hasselbalch, buffer, negative pH, polyprotic acid.