Titration Calculator

Calculate unknown concentration from acid-base titration data

Titration Data

Known concentration (M) of titrant
Volume (mL) used to reach endpoint
Volume (mL) of unknown solution
H+ (acid) or OH- (base) per titrant formula unit, e.g. HCl 1, NaOH 1, H2SO4 2
H+ (acid) or OH- (base) per analyte formula unit, e.g. H2SO4 2, Ca(OH)2 2

Results

Analyte Concentration
0.00 M
Unknown solution concentration

Understanding Acid-Base Titration

Titration is a fundamental analytical technique in chemistry used to determine the concentration of an unknown solution by reacting it with a solution of known concentration. Acid-base titration, specifically, involves neutralization reactions between acids and bases, and is one of the most common types of volumetric analysis.

What is Titration?

In a titration, a solution of known concentration (the titrant) is gradually added to a known volume of solution with unknown concentration (the analyte) until the reaction reaches completion. This completion point, called the equivalence point, is often detected using an indicator that changes color or by monitoring pH.

n₁M₁V₁ = n₂M₂V₂

Where n is the n-factor (number of H+ or OH- ions per molecule).

Why Titration Matters

  • Quantitative analysis: Accurately determine unknown concentrations
  • Quality control: Verify product concentrations in manufacturing
  • Environmental testing: Measure water quality and pollutant levels
  • Food industry: Determine acidity in beverages and foods
  • Pharmaceuticals: Assay drug concentrations
  • Research: Characterize new compounds and reactions

Titration Calculation Formula

For simple 1:1 stoichiometry (like HCl + NaOH):

M₁V₁ = M₂V₂

For reactions with different stoichiometry (like H₂SO₄ + 2NaOH):

n₁M₁V₁ = n₂M₂V₂

Where n represents the n-factor or number of acidic hydrogens (for acids) or hydroxide ions (for bases).

Example Calculation

25.0 mL of 0.1 M HCl is used to titrate 20.0 mL of NaOH solution. Find the concentration of NaOH:

  • M₁V₁ = M₂V₂
  • (0.1 M)(25.0 mL) = M₂(20.0 mL)
  • M₂ = (0.1 × 25.0) / 20.0
  • M₂ = 0.125 M

Types of Acid-Base Titrations

  • Strong acid - strong base: Sharp pH change at equivalence point (pH 7)
  • Strong acid - weak base: Equivalence point pH < 7
  • Weak acid - strong base: Equivalence point pH > 7
  • Weak acid - weak base: No sharp endpoint, rarely used

Common Indicators

  • Phenolphthalein: Colorless to pink (pH 8.2-10), used for strong acid-base titrations
  • Methyl orange: Red to yellow (pH 3.1-4.4), used for strong acid-weak base
  • Bromothymol blue: Yellow to blue (pH 6.0-7.6), general purpose
  • Universal indicator: Multiple color changes across pH range

Equipment and Procedure

Essential titration equipment includes:

  • Burette: For delivering precise volumes of titrant
  • Pipette: For measuring exact volumes of analyte
  • Erlenmeyer flask: For mixing the solutions
  • pH meter or indicator: For detecting the endpoint
  • White tile: To better see color changes

Standard Solutions

The titrant must be a standard solution - one with accurately known concentration. Primary standards are highly pure compounds used to prepare standard solutions. Common primary standards include:

  • Sodium carbonate (Na₂CO₃) - for standardizing acids
  • Potassium hydrogen phthalate (KHP) - for standardizing bases
  • Oxalic acid (H₂C₂O₄) - for standardizing bases

Sources of Error

  • Parallax error: Reading burette from wrong angle
  • Air bubbles: In burette tip affecting volume
  • Overshooting endpoint: Adding titrant too quickly
  • Contamination: Dirty glassware affecting results
  • Incorrect indicator: Choosing wrong pH range
  • Temperature effects: Volume changes with temperature

Tips for Accurate Titrations

  • Rinse burette with titrant solution before filling
  • Remove air bubbles from burette tip
  • Add titrant slowly near the endpoint
  • Swirl flask constantly during titration
  • Perform multiple trials and average results
  • Read meniscus at eye level
  • Use appropriate indicator for the titration type

Beyond Acid-Base Titrations

Other types of titrations include:

  • Redox titrations: Involve electron transfer (e.g., KMnO₄ titrations)
  • Complexometric titrations: Form complex ions (e.g., EDTA for metal ions)
  • Precipitation titrations: Form precipitates (e.g., silver nitrate for chloride)

Applications in Industry

  • Water treatment: Measuring alkalinity and hardness
  • Wine making: Determining titratable acidity
  • Dairy industry: Testing milk acidity
  • Chemical manufacturing: Quality control of products
  • Petroleum industry: Measuring acid or base number of oils

Whether you're a chemistry student performing laboratory experiments, an analyst conducting quality control tests, or a researcher quantifying compounds, our titration calculator provides quick, accurate concentration determinations from your titration data.

Frequently Asked Questions

Worked example: 18.4 mL of 0.1 M HCl neutralizes 25.0 mL of Ca(OH)2. What is the Ca(OH)2 concentration?

Enter HCl as the titrant (titrant n-factor 1) and Ca(OH)2 as the analyte (analyte n-factor 2). From n_analyte × C_analyte × V_analyte = n_titrant × C_titrant × V_titrant: 2 × C × 25.0 = 1 × 0.1 × 18.4, so C = 1.84 ÷ 50 = 0.0368 M. As a check, 0.1 × 0.0184 L = 0.00184 mol HCl neutralizes 0.00092 mol Ca(OH)2, and 0.00092 ÷ 0.025 L = 0.0368 M.

What equation does this titration calculator solve?

It uses n_analyte × C_analyte × V_analyte = n_titrant × C_titrant × V_titrant, so C_analyte = (n_titrant × C_titrant × V_titrant) ÷ (n_analyte × V_analyte). Each n-factor counts the H+ an acid donates (HCl 1, H2SO4 2) or the OH− a base supplies (NaOH 1, Ca(OH)2 2), whichever side the acid or base is on. Both volumes are in mL; the units cancel. The result is in mol/L and also shown in mM.

How do I handle a base titrant, such as NaOH titrating sulfuric acid?

Enter the n-factor of each substance in its own box. If 30.0 mL of 0.1 M NaOH neutralizes 25.0 mL of H2SO4, the titrant n-factor is 1 (NaOH) and the analyte n-factor is 2 (H2SO4). C = (1 × 0.1 × 30.0) ÷ (2 × 25.0) = 0.0600 M. That matches the mole check: 0.003 mol NaOH reacts with 0.0015 mol H2SO4, which is 0.0600 M in 0.025 L.

What is the difference between the endpoint and the equivalence point?

The equivalence point is where the moles of H+ and OH− supplied are exactly equal, which is the condition the formula assumes. The endpoint is where the indicator changes color, and it should fall close to the equivalence point. Choose an indicator whose color-change range brackets the equivalence pH. Phenolphthalein (pH 8.2–10) suits weak acid–strong base titrations, and methyl orange (3.1–4.4) suits strong acid–weak base.