Experiment: Determining pH of Common Solutions - Tenth Grade Science

The hydrogen ion concentration of an aqueous solution governs its chemical reactivity, physiological compatibility, and industrial utility. In this laboratory investigation, students use universal indicator paper and standardized pH colorimetric charts to experimentally measure the pH values of common household and laboratory solutions, categorizing them systematically into acidic, neutral, or basic classifications.

Experimental Aim and Objectives

The primary objectives of this practical investigation are:
1. To determine the pH values of various test samples—including dilute hydrochloric acid, dilute sodium hydroxide solution, ethanoic acid, lemon juice, pure distilled water, and sodium hydrogen carbonate solution—using universal indicator paper and broad-range pH indicator strips.

2. To correlate observed indicator color changes with numeric pH values on the standardized 0 to 14 pH scale.

3. To classify each tested liquid sample as a strong acid, weak acid, neutral solution, weak base, or strong base based on measured hydronium ion concentration.


Theoretical Background and Chemical Principles

The pH scale was introduced in 1909 by the Danish biochemist Søren Peder Lauritz Sørensen to quantify the molar concentration of hydrogen ions ($[H^+]$ or $[H_3O^+]$) in aqueous solutions.

                  THE STANDARDIZED pH SPECTRUM

  pH Value:   0   1   2   3   4   5   6   7   8   9  10  11  12  13  14
  Color:     RED      ORANGE     YELLOW  GREEN   BLUE      VIOLET  PURPLE
  Nature:    <--- STRONGLY ACIDIC --->   NEUTRAL   <--- STRONGLY BASIC --->

Mathematical Formulation

The pH of an aqueous solution is defined mathematically as the negative common (base-10) logarithm of the active molar hydrogen ion concentration:

$$\text{pH} = -\log_{10}[H_3O^+] \quad \text{or} \quad [H_3O^+] = 10^{-\text{pH}}$$

In pure neutral water at $25^\circ\text{C}$, auto-ionization yields equal concentrations of hydronium and hydroxide ions ($[H_3O^+] = [OH^-] = 1.0 \times 10^{-7} \text{ mol/L}$), producing a neutral pH of 7.0:

$$\text{pH} = -\log_{10}(1.0 \times 10^{-7}) = 7.0$$

Solutions with $[H_3O^+] > 1.0 \times 10^{-7} \text{ mol/L}$ yield $\text{pH} < 7$ and are classified as acidic. Solutions with $[H_3O^+] < 1.0 \times 10^{-7} \text{ mol/L}$ yield $\text{pH} > 7$ and are classified as alkaline (basic).

Universal Indicator Chemistry

Unlike single-threshold indicators (such as litmus, which simply toggles red or blue), a universal indicator is a balanced blend of multiple synthetic indicator dyes—including thymol blue, methyl red, bromothymol blue, and phenolphthalein. As hydronium ion activity changes, the dyes undergo sequential protonation and deprotonation equilibria, producing a smooth gradient of color transitions across the full 0 to 14 pH spectrum: Strongly Acidic ($\text{pH } 0 \text{ to } 2$): Deep scarlet red. Moderately Acidic ($\text{pH } 3 \text{ to } 5$): Orange to reddish-yellow. Weakly Acidic ($\text{pH } 6$): Pale greenish-yellow. Neutral ($\text{pH } 7$): Pure emerald green. Weakly Basic ($\text{pH } 8 \text{ to } 9$): Cyan to grayish-blue. Moderately Basic ($\text{pH } 10 \text{ to } 11$): Deep navy blue. Strongly Basic ($\text{pH } 12 \text{ to } 14$): Violet to deep purple.

Apparatus and Reagents Required

                  LABORATORY BENCHTOP ARRANGEMENT

      +-----------------------------------------------------------+
      |  [ Test-Tube Rack ]          [ Ceramic Spotting Plate ]   |
      |   1  2  3  4  5  6                (o) (o) (o)             |
      |  ( )( )( )( )( )( )               (o) (o) (o)             |
      |                                                           |
      |  [ Dropper Bottles ]         [ Universal Indicator Strips]|
      |  [ Watch Glasses   ]         [ Standardized Color Chart ] |
      +-----------------------------------------------------------+

Chemical Reagents

Dilute Hydrochloric acid ($0.1\text{ M } HCl$). Dilute Sodium hydroxide solution ($0.1\text{ M } NaOH$). Dilute Ethanoic acid (acetic acid, $CH_3COOH$). Freshly extracted lemon juice (diluted 1:1 with distilled water). Pure distilled water ($H_2O$, deionized). Sodium hydrogen carbonate solution ($0.1\text{ M } NaHCO_3$).

Laboratory Apparatus

Clean porcelain white ceramic spotting tile (with multiple sample depressions). High-precision universal indicator paper strip booklet with calibrated color reference chart. Six clean, dry glass droppers or glass stirring rods. Six labeled $50\text{ mL}$ borosilicate beakers. Deionized wash bottle with distilled water. Clean absorbent filter paper strips.

Step-by-Step Experimental Procedure

  1. Wash the ceramic spotting tile thoroughly with distilled water and dry it with clean absorbent paper. Label the depressions corresponding to the six test solutions.
  2. Pour 10 to 15 mL of each test liquid into its labeled beaker.
  3. Tear six clean strips of universal indicator paper (each roughly 2 cm in length) using clean tweezers, and place one strip into each dry depression of the white ceramic spotting tile.
  4. Using a clean glass dropper, withdraw a few drops of the first solution (dilute hydrochloric acid) and deposit two drops directly onto the first indicator paper strip.
  5. Observe the immediate color change that develops across the moistened strip against the white porcelain background.
  6. Compare the developed color with the standardized manufacturer color reference chart, matching the shade to record the corresponding numeric pH value.
  7. Rinse the glass dropper thoroughly with distilled water before moving to the next sample, or use a fresh, dry dropper for each test liquid.
  8. Repeat steps 4 through 7 for the remaining samples: dilute sodium hydroxide, ethanoic acid, lemon juice, distilled water, and sodium hydrogen carbonate solution.
  9. Record all observed colors and corresponding pH values in the experimental observation table.
               SAMPLING ON PORCELAIN SPOTTING TILE

        Sample Depressions:
        +-----------------------------------------------+
        |  (1) HCl         (2) CH3COOH      (3) Lemon   |
        |   [ Red ]          [ Orange ]      [ Pink-Red]|
        |                                               |
        |  (4) Dist. Water (5) NaHCO3       (6) NaOH    |
        |   [ Green ]        [ Blue ]        [ Purple ] |
        +-----------------------------------------------+

Observations and Experimental Results

                  EXPERIMENTAL OBSERVATION TABLE

  Sample  Test Solution         Observed Color     Inferred pH  Classification
  ----------------------------------------------------------------------------
  1       Dilute HCl (0.1 M)    Dark Red           1.0          Strong Acid
  2       Ethanoic Acid (0.1 M) Orange             3.0 to 4.0   Weak Acid
  3       Fresh Lemon Juice     Pinkish-Red        2.0 to 2.5   Moderate Acid
  4       Distilled Water       Grass Green        7.0          Neutral
  5       Dilute NaHCO3 (0.1 M) Light Blue         8.5 to 9.0   Weak Base
  6       Dilute NaOH (0.1 M)   Deep Violet-Purple 13.0 to 14.0 Strong Base

Chemical Explanations and Ionic Dissociations

  1. Hydrochloric Acid: A strong mineral acid that dissociates completely in water ($HCl + H_2O \to H_3O^+ + Cl^-$). A $0.1\text{ M}$ solution yields $[H_3O^+] \approx 10^{-1} \text{ mol/L}$, resulting in a low pH of 1.0 and a dark red color.
  2. Ethanoic Acid: A weak organic carboxylic acid that undergoes only partial dissociation in water ($CH_3COOH + H_2O \rightleftharpoons H_3O^+ + CH_3COO^-$), with an ionization constant $K_a \approx 1.8 \times 10^{-5}$. A $0.1\text{ M}$ solution yields a modest $[H_3O^+] \approx 1.34 \times 10^{-3} \text{ mol/L}$, producing a pH near 3.0 and an orange color.
  3. Lemon Juice: Contains naturally occurring citric acid ($C_6H_8O_7$), a tricarboxylic acid that produces an acidic pH around 2.2, turning the indicator paper pinkish-red.
  4. Distilled Water: High-purity deionized water contains equal concentrations of hydronium and hydroxide ions ($[H_3O^+] = [OH^-] = 10^{-7} \text{ mol/L}$), yielding a neutral pH of 7.0 and a green indicator response.
  5. Sodium Hydrogen Carbonate: An amphoteric salt that undergoes alkaline hydrolysis in water ($HCO_3^- + H_2O \rightleftharpoons H_2CO_3 + OH^-$), generating excess hydroxide ions. This weak base yields a pH around 8.5 to 9.0, turning the indicator paper light blue.
  6. Sodium Hydroxide: A strong alkali that dissociates completely into ions ($NaOH \to Na^+ + OH^-$). A $0.1\text{ M}$ solution yields $[OH^-] = 10^{-1} \text{ mol/L}$. From the ion-product of water ($K_w = [H_3O^+][OH^-] = 10^{-14}$), $[H_3O^+] = 10^{-13} \text{ mol/L}$, giving a high pH of 13.0 and a deep violet-purple color.

Laboratory Precautions and Sources of Error

  1. Dropper Cross-Contamination: Always rinse glass droppers thoroughly with distilled water between tests, or use separate dedicated droppers for each reagent. A single drop of residual concentrated sodium hydroxide will neutralize an acidic test sample, skewing the measured pH.
  2. Direct Dipping Hazard: Never insert indicator paper strips directly into master reagent bottles. Always transfer small aliquots of test solutions to a spotting tile or watch glass to avoid contaminating the main reagent supply.
  3. Rapid Color Reading: Record color matches within 10 to 15 seconds of moistening the indicator strip. Allowing the strip to dry in air exposes it to atmospheric carbon dioxide, which forms carbonic acid ($H_2CO_3$) and slowly shifts alkaline indicator colors toward acidic hues.
  4. Optical Reading Conditions: Evaluate color matches under natural daylight or bright, balanced white laboratory illumination. Low ambient light or yellow incandescent bulbs can distort color perception, leading to misreadings.

Chapter Exercises and Solutions

Exercise 1: A student tests an unknown clear liquid using universal indicator paper and observes that the paper turns a pale greenish-yellow color, matching a pH value of 6.0 on the reference chart. Calculate the molar hydronium ion concentration $[H_3O^+]$ of this unknown sample, and classify its chemical nature.

Solution 1: The mathematical relationship between pH and hydronium ion concentration is given by the definition $\text{pH} = -\log_{10}[H_3O^+]$, which rearranges to $[H_3O^+] = 10^{-\text{pH}} \text{ mol/L}$. For a measured pH value of 6.0, the hydronium ion concentration is $[H_3O^+] = 10^{-6.0} = 1.0 \times 10^{-6} \text{ mol/L}$. Because this hydronium concentration is greater than the neutral threshold ($1.0 \times 10^{-7} \text{ mol/L}$) and the pH is less than 7.0, the liquid is classified as a weakly acidic solution (such as fresh milk or unpolluted rainwater containing dissolved atmospheric carbon dioxide).

Exercise 2: Explain why a $0.1\text{ M}$ solution of hydrochloric acid yields a pH of approximately 1.0, whereas a $0.1\text{ M}$ solution of ethanoic acid yields a pH of approximately 3.0, despite both solutions possessing identical molar concentrations of acid molecules.

Solution 2: The measured pH reflects the equilibrium concentration of free hydronium ions in solution, rather than the total quantity of dissolved acid molecules. Hydrochloric acid ($HCl$) is a strong mineral acid that undergoes complete ionization in water: $HCl + H_2O \to H_3O^+ + Cl^-$. In a $0.1\text{ M } HCl$ solution, every dissolved molecule dissociates, yielding $[H_3O^+] = 0.1\text{ M} = 10^{-1} \text{ mol/L}$, which gives $\text{pH} = -\log_{10}(10^{-1}) = 1.0$. In contrast, ethanoic acid ($CH_3COOH$) is a weak carboxylic acid that ionizes only partially in water, establishing a dynamic chemical equilibrium: $CH_3COOH + H_2O \rightleftharpoons H_3O^+ + CH_3COO^-$. With an ionization constant $K_a \approx 1.8 \times 10^{-5}$, only about 1.3 percent of ethanoic acid molecules dissociate in a $0.1\text{ M}$ solution, yielding $[H_3O^+] \approx 1.34 \times 10^{-3} \text{ mol/L}$. This lower hydronium ion concentration yields $\text{pH} = -\log_{10}(1.34 \times 10^{-3}) \approx 2.87 \approx 3.0$.

Exercise 3: If a strip of universal indicator paper moistened with a basic sodium hydroxide solution is left exposed on an open laboratory bench for thirty minutes, describe what color change will occur over time, explaining the underlying chemical reactions with ambient air.

Solution 3: When moistened with sodium hydroxide solution, the universal indicator paper initially displays a deep violet-purple color, indicating a strongly alkaline pH between 13.0 and 14.0. If the strip is left exposed to ambient air for thirty minutes, the indicator color will shift from purple through blue and cyan, eventually turning green or yellowish-green (corresponding to a pH of 7.0 to 6.5). This shift occurs because atmospheric air contains carbon dioxide gas ($CO_2$). The carbon dioxide dissolves into the moisture on the strip, reacting with the alkaline sodium hydroxide to form sodium carbonate and water: $2 NaOH + CO_2 \to Na_2CO_3 + H_2O$. Continued absorption of carbon dioxide converts the carbonate into sodium hydrogen carbonate: $Na_2CO_3 + CO_2 + H_2O \to 2 NaHCO_3$. This progressive neutralization depletes free hydroxide ions ($OH^-$), lowering the solution pH and shifting the indicator color across the spectrum.