# pH – Acids & Buffers ## Overview - PAG 11.3 investigates pH changes in different types of solutions through three distinct parts: - Preparing hydrochloric acid solutions at different concentrations by dilution and measuring their pH - Preparing buffer solutions and measuring their pH - Investigating how pH changes when sodium hydroxide solution is added to both acid and buffer solutions

This practical develops quantitative skills in acid-base chemistry and demonstrates why buffers are so important in biological and industrial systems

## Part 1: Preparing Hydrochloric Acid Solutions ### Serial Dilution Technique - Serial dilution systematically reduces concentration by the same factor at each step - Each dilution creates a solution that's a fixed fraction of the previous concentration **Common dilution ratios:** | Starting Volume | Water Added | Total Volume | Dilution Ratio | |----------------|-------------|--------------|----------------| | 1.0 cm³ | 9.0 cm³ | 10.0 cm³ | 1:10 | | 1.0 cm³ | 4.0 cm³ | 5.0 cm³ | 1:5 | | 2.0 cm³ | 8.0 cm³ | 10.0 cm³ | 1:5 |

A **1:10 dilution** means the concentration decreases to one-tenth of the original value at each step

### Method 1. Label 7 test tubes numbered 1 to 7 2. Add 10.0 cm³ of 0.1 mol dm⁻³ HCl to test tube 1 - Test tube 1 is undiluted - Contains 0.1 mol dm⁻³ HCl 3. Transfer 1.0 cm³ from test tube 1 into a 10.0 cm³ measuring cylinder 4. Add 9.0 cm³ of distilled water to the measuring cylinder - Total volume = 10.0 cm³ 5. Transfer this solution to test tube 2 - This is a 1:10 dilution - Test tube 2 contains 0.01 mol dm⁻³ HCl 6. Repeat steps 3-5: - Use test tube 2 as the source for test tube 3 - Continue the pattern through to test tube 6 7. Test tube 7 should contain only distilled water (blank control) 8. Measure the pH of each solution using: - A calibrated pH probe, OR - Universal indicator with a colour comparison chart ### Specimen Results for Serial Dilution - Strong acids show a predictable pH pattern during serial dilution - Each 10-fold decrease in concentration increases pH by approximately 1 unit | Test Tube | Concentration (mol dm⁻³) | pH | |-----------|-------------------------|-----| | 1 | 0.1 | 1.00 | | 2 | 0.01 | 2.00 | | 3 | 0.001 | 3.00 | | 4 | 0.0001 | 4.00 | | 5 | 0.00001 | 5.00 | | 6 | 0.000001 | 6.00 | | 7 | 0.0000001 | 7.00 | **Visualising the logarithmic relationship:** The graph below shows how pH increases linearly as concentration decreases exponentially. This demonstrates the fundamental definition of pH as the negative logarithm of hydrogen ion concentration.

The linear relationship between $\log[\text{H}^+]$ and pH demonstrates that pH is defined as $\text{pH} = -\log[\text{H}^+]$

### Testing with Sodium Hydroxide - These acid solutions can be titrated with NaOH to monitor pH change throughout neutralisation - Use a calibrated pH meter to track pH continuously during titration - The resulting graphs should show characteristic strong acid-strong base titration curves with steep vertical sections at the equivalence point ## Part 2: Preparing Buffer Solutions

A **buffer solution** resists pH changes when small amounts of acid or base are added

### Understanding Buffer Capacity - Buffer resistance to pH change has a finite limit called **buffer capacity** - Higher concentrations of buffer components = greater buffer capacity - This practical prepares two buffers with different capacities for comparison ### Method for Standard Buffer Solution **Initial calculations:** - Starting solution: 5 cm³ of 0.30 mol dm⁻³ ethanoic acid - Final volume: 100 cm³ - Calculate the concentration after dilution: $$\text{moles of ethanoic acid} = \frac{0.3}{1000} \times 5 = 0.0015 \text{ moles}$$ $$\text{concentration in 100 cm}^3 = \frac{0.0015}{100} \times 1000 = 0.015 \text{ mol dm}^{-3}$$ **Calculating sodium ethanoate mass:** - Use the Henderson-Hasselbalch equation: $$\text{pH} = \text{pK}_\text{a} + \log_{10}\frac{[\text{A}^-]}{[\text{HA}]}$$ - For ethanoic acid: $\text{pK}_\text{a} = 4.76$ - Target pH = 5.00 $$5.00 = 4.76 + \log_{10}\frac{[\text{A}^-]}{0.015}$$ $$0.24 = \log_{10}\frac{[\text{A}^-]}{0.015}$$ $$10^{0.24} = \frac{[\text{A}^-]}{0.015}$$ $$[\text{A}^-] = 0.0261 \text{ mol dm}^{-3}$$ - Moles of A⁻ needed in 100 cm³: $$\frac{0.0261}{10} = 0.00261 \text{ moles}$$ - Mass of sodium ethanoate required (Mr = 82.04): $$0.00261 \times 82.04 = 0.214 \text{ g}$$ **Practical steps:** 1. Add 50 cm³ of distilled water to a 100 cm³ beaker 2. Add 5 cm³ of the ethanoic acid using a pipette 3. Measure approximately 0.22 g of sodium ethanoate into a weighing boat - This slightly higher mass accounts for transfer loss - Some material will remain in the weighing boat 4. Calibrate a pH meter and place the probe in the beaker 5. Slowly stir in the sodium ethanoate with a glass rod - **Never use the pH probe as a stirring rod** 6. Stop adding sodium ethanoate when pH reaches 5.00 7. Reweigh the sodium ethanoate boat to determine actual mass used 8. Transfer the buffer solution to a 100 cm³ volumetric flask 9. Dilute to the mark with distilled water, cap and mix thoroughly

Always stir buffer solutions with a glass rod, not the pH probe—probes are fragile and easily damaged

### Method for Higher Capacity Buffer **Why make a higher capacity buffer?** - Comparing buffers with different capacities demonstrates the relationship between concentration and buffer effectiveness - Higher concentration buffers can neutralise more added acid or base before pH changes significantly **Modified calculations:** - Starting solution: 5 cm³ of **0.50 mol dm⁻³** ethanoic acid (instead of 0.30 mol dm⁻³) - Method remains identical to the standard buffer - Only the calculations change: $$\text{moles of ethanoic acid} = \frac{0.5}{1000} \times 5 = 0.0025 \text{ moles}$$ $$\text{concentration in 100 cm}^3 = \frac{0.0025}{100} \times 1000 = 0.025 \text{ mol dm}^{-3}$$ - Using the same target pH of 5.00: $$5.00 = 4.76 + \log_{10}\frac{[\text{A}^-]}{0.025}$$ $$[\text{A}^-] = 0.0434 \text{ mol dm}^{-3}$$ - Moles needed in 100 cm³: $$\frac{0.0434}{10} = 0.00434 \text{ moles}$$ - Mass of sodium ethanoate required: $$0.00434 \times 82.04 = 0.356 \text{ g}$$ ## Part 3: Testing Buffer Capacity ### pH Probe Calibration

Always calibrate a pH probe using a **two-point calibration** with pH 4 and pH 9 buffer standards before use

**Calibration procedure:** 1. Rinse the pH probe with distilled water 2. Gently shake to remove excess water (don't wipe—this can cause static) 3. Place the tip of the probe in pH 4 buffer solution - Ensure the glass bulb is fully immersed 4. Allow the reading to stabilise 5. Adjust the reading to exactly 4.0 6. Rinse with distilled water and shake dry 7. Place the tip in pH 9 buffer solution 8. Check the reading—it should be very close to 9.0 9. If adjustment is needed, recalibrate with pH 4 first, then repeat

pH probes lose accuracy during storage as the glass electrode becomes unstable—regular calibration is essential for reliable results

### Comparing Buffer Capacities **Method:** 1. Calibrate the pH probe using pH 4 and 9 buffer standards 2. Rinse with distilled water before and between all measurements 3. Allow readings to stabilise before recording values 4. Transfer 25.0 cm³ of the standard buffer to a conical flask using a volumetric pipette 5. Fill a burette with standardised NaOH solution 6. Monitor pH continuously using the calibrated probe 7. Add NaOH from the burette in small increments 8. Record the volume of NaOH needed to cause a 1-unit pH increase 9. Repeat the entire process with 25.0 cm³ of the higher-concentration buffer 10. Compare results to evaluate which buffer has greater capacity **What to look for:** - The higher-concentration buffer should require **more NaOH** to achieve the same pH change - This demonstrates that buffer capacity increases with concentration - Real-world applications: biological systems use highly concentrated buffers to maintain stable pH despite metabolic acid/base production ### Analysis and Interpretation **Expected observations:** | Buffer Type | Approximate Volume NaOH for +1 pH unit | Buffer Capacity | |-------------|----------------------------------------|-----------------| | 0.015 mol dm⁻³ | ~5-8 cm³ | Lower | | 0.025 mol dm⁻³ | ~8-13 cm³ | Higher | - The buffer with higher concentrations of both weak acid and conjugate base can neutralise more added base - Eventually, all buffers reach their capacity limit and pH begins to rise steeply - This occurs when one component of the conjugate acid-base pair is exhausted **Visualising buffer action:** The graph below shows specimen data comparing how pH changes when NaOH is added to a buffer solution versus an unbuffered acid solution. Notice how the buffer maintains relatively stable pH over a much larger volume of added base.

Think of buffer capacity like a sponge—a thicker sponge (higher concentration) can absorb more water (added acid/base) before becoming saturated

## Key Practical Skills Summary **This practical develops essential techniques:** - Performing accurate serial dilutions to create a concentration range - Using pH meters and indicators for pH measurement - Calibrating pH probes correctly using two-point calibration - Preparing buffer solutions from calculated quantities of weak acid and conjugate base - Using the Henderson-Hasselbalch equation to design a buffer: $\text{pH} = \text{pK}_\text{a} + \log_{10}\frac{[\text{A}^-]}{[\text{HA}]}$ - Monitoring pH changes during titration - Comparing experimental results to evaluate buffer capacity **Understanding developed:** - How serial dilution affects pH in predictable ways for strong acids - Why buffers resist pH change through conjugate acid-base pairs - The relationship between buffer concentration and buffer capacity - How to select appropriate buffer concentrations for practical applications - The limitations of buffer action when capacity is exceeded **Common sources of error:** - Inaccurate dilution volumes lead to incorrect concentrations - Uncalibrated or poorly calibrated pH probes give unreliable readings - Transfer losses when adding solids (sodium ethanoate) - Contamination from previous solutions if glassware isn't properly rinsed - Stirring with the pH probe can damage the fragile glass electrode

Always check pH probe calibration at both pH 4 and pH 9—if the probe doesn't read close to 9.0 after calibrating at 4.0, the electrode may need replacing

## Connecting to Theory **Strong acid pH behaviour:** - For a strong monobasic acid: $[\text{H}^+] = [\text{acid}]$ - Therefore: $\text{pH} = -\log[\text{acid}]$ - Each 10-fold dilution increases pH by exactly 1 unit - This linear relationship confirms complete dissociation **Buffer action:** - A buffer contains both a weak acid (HA) and its conjugate base (A⁻) - When H⁺ is added: the conjugate base neutralises it → $\text{A}^- + \text{H}^+ \rightarrow \text{HA}$ - When OH⁻ is added: the weak acid neutralises it → $\text{HA} + \text{OH}^- \rightarrow \text{A}^- + \text{H}_2\text{O}$ - Both reactions shift equilibria but cause minimal pH change - Buffer effectiveness depends on having significant concentrations of both species **Real-world relevance:** - Blood plasma uses the carbonic acid-hydrogencarbonate buffer system: $\text{H}_2\text{CO}_3 / \text{HCO}_3^-$ - Maintains blood pH between 7.35 and 7.45 despite metabolic CO₂ production - pH outside this narrow range causes serious medical problems (acidosis or alkalosis) - Many industrial processes require pH control through buffer systems - Fermentation, pharmaceutical production, and water treatment all rely on buffers