What is pH?
pH is a measure of the acidity or basicity of an aqueous solution. The term "pH" stands for "power of hydrogen" and represents the negative logarithm (base 10) of the hydrogen ion concentration. The pH scale typically ranges from 0 to 14, though values outside this range are possible for very concentrated solutions.
Understanding pH is crucial in chemistry, biology, environmental science, medicine, and many industrial processes. From maintaining proper blood pH to treating wastewater, pH measurements play a vital role in our daily lives.
The pH Scale Explained
Acidic Solutions (pH 0-6.9)
Solutions with pH less than 7 are acidic. The lower the pH, the higher the concentration of hydrogen ions [H⁺]. Strong acids like hydrochloric acid (HCl) and sulfuric acid (H₂SO₄) have pH values close to 0. Weak acids like acetic acid (vinegar) typically have pH around 2-3.
Neutral Solutions (pH 7)
Pure water at 25°C (77°F) has a pH of exactly 7.0, meaning it has equal concentrations of H⁺ and OH⁻ ions (each 1.0 × 10⁻⁷ M). At this point, the solution is neither acidic nor basic. However, neutral pH can vary slightly with temperature.
Basic/Alkaline Solutions (pH 7.1-14)
Solutions with pH greater than 7 are basic or alkaline. These solutions have higher concentrations of hydroxide ions [OH⁻] than hydrogen ions [H⁺]. Strong bases like sodium hydroxide (NaOH) can have pH values approaching 14.
pH Calculations and Formulas
Calculating pH from [H⁺]
The fundamental formula for pH is: pH = -log₁₀[H⁺], where [H⁺] is the hydrogen ion concentration in moles per liter (M). For example, if [H⁺] = 1.0 × 10⁻⁴ M, then pH = -log(1.0 × 10⁻⁴) = 4.0.
Calculating [H⁺] from pH
To find hydrogen ion concentration from pH, use: [H⁺] = 10⁻ᵖᴴ. For pH 3, [H⁺] = 10⁻³ = 0.001 M = 1.0 × 10⁻³ M.
The pOH Scale
Similar to pH, pOH measures hydroxide ion concentration: pOH = -log₁₀[OH⁻]. At 25°C, pH and pOH are related by: pH + pOH = 14. This relationship comes from the ion product of water (Kw).
Ion Product of Water (Kw)
At 25°C, the product of hydrogen and hydroxide ion concentrations is constant: [H⁺] × [OH⁻] = 1.0 × 10⁻¹⁴. This equilibrium constant is called Kw and is fundamental to all pH calculations.
How to Use Our pH Calculator
Calculate from pH
Enter a pH value between 0 and 14. The calculator instantly computes pOH, [H⁺], and [OH⁻]. The visual pH scale shows where your solution falls on the acidity-basicity spectrum with color coding.
Calculate from pOH
Input pOH value, and the calculator determines pH, [H⁺], and [OH⁻]. This is useful when you know the base concentration and want to find acidity measures.
Calculate from [H⁺]
Enter hydrogen ion concentration in mol/L using scientific notation (1e-7) or decimal format. The calculator computes all related values including pH, pOH, and [OH⁻].
Calculate from [OH⁻]
Input hydroxide ion concentration to find pH, pOH, and [H⁺]. Useful for basic solutions where OH⁻ concentration is known.
Common pH Values
Strong Acids (pH 0-2)
Battery acid (pH ~0): Highly corrosive sulfuric acid. Gastric acid (pH ~1-2): Stomach acid for digestion. Lemon juice (pH ~2): Contains citric acid. These substances require careful handling.
Weak Acids (pH 3-6)
Vinegar (pH ~2.5): Acetic acid solution. Orange juice (pH ~3-4): Contains citric acid. Black coffee (pH ~5): Mildly acidic. Milk (pH ~6): Slightly acidic.
Neutral (pH ~7)
Pure water (pH 7.0): Perfectly neutral at 25°C. Blood (pH 7.35-7.45): Slightly basic, tightly regulated. Saliva (pH 6.5-7.5): Varies with diet and hydration.
Weak Bases (pH 8-10)
Seawater (pH ~8.1): Slightly basic. Baking soda (pH ~8.5): Sodium bicarbonate solution. Milk of magnesia (pH ~10): Antacid suspension.
Strong Bases (pH 11-14)
Ammonia (pH ~11): Common cleaning agent. Bleach (pH ~12-13): Sodium hypochlorite solution. Drain cleaner (pH ~14): Contains sodium hydroxide. Extremely caustic.
Applications of pH Measurements
Human Health
Blood pH must remain between 7.35-7.45 for proper physiological function. Acidosis (pH < 7.35) or alkalosis (pH > 7.45) can be life-threatening. Stomach pH (~1-2) is necessary for digestion but can cause ulcers if not properly regulated.
Agriculture and Soil Science
Soil pH (typically 4-9) affects nutrient availability for plants. Most crops prefer slightly acidic to neutral soil (pH 6-7). Soil pH testing guides fertilizer and amendment choices to optimize plant growth.
Aquariums and Aquatic Life
Freshwater fish typically require pH 6.5-7.5, while marine fish need pH 8.1-8.4. Maintaining stable pH is crucial for fish health. Ammonia toxicity increases dramatically at higher pH values.
Swimming Pools
Pool water should maintain pH 7.2-7.8 for comfort and effective chlorine disinfection. Low pH causes eye irritation and equipment corrosion. High pH reduces chlorine effectiveness and causes cloudiness.
Water Treatment
Municipal water treatment requires pH adjustment for disinfection efficiency and pipe protection. Drinking water typically has pH 6.5-8.5. Wastewater treatment uses pH control for biological processes and pollutant removal.
Food and Beverage Industry
pH affects food safety, flavor, texture, and shelf life. Acidic foods (pH < 4.6) resist bacterial growth. pH monitoring ensures product quality and prevents spoilage in fermentation, canning, and preservation processes.
Chemical Manufacturing
Many industrial processes require precise pH control. Chemical reactions, crystallization, and purification often depend on specific pH ranges. pH monitoring prevents equipment corrosion and ensures product specifications.
pH Measurement Methods
pH Paper (Litmus Paper)
Simple color-change indicators provide approximate pH (±1 unit). Blue litmus turns red in acid; red litmus turns blue in base. Inexpensive and convenient but less accurate than other methods.
Universal Indicator
Multi-component indicator that displays different colors across the pH range. Provides reasonably accurate pH readings (±0.5 units) by comparing color to a chart. Good for educational purposes and field testing.
pH Meter (Electrode)
Electronic device measuring pH precisely (±0.01 units). Uses glass electrode sensitive to H⁺ ions. Requires calibration with buffer solutions. Industry standard for accurate pH determination in laboratories and industrial settings.
pH Test Strips
Pre-treated paper strips with multiple indicator zones. Quick and easy to use. Suitable for routine testing of pools, aquariums, and body fluids. Accuracy typically ±0.5 pH units.
Factors Affecting pH
Temperature
pH varies with temperature because water's ion product (Kw) changes. Pure water has pH 7.0 at 25°C but pH 7.5 at 0°C and pH 6.5 at 60°C. Always specify temperature when reporting pH values.
Dilution
Diluting acids or bases changes their pH. However, dilution has less effect near pH 7 due to water's buffering capacity. Strong acids and bases show dramatic pH changes when diluted.
Buffer Solutions
Buffers resist pH changes when small amounts of acid or base are added. Consist of weak acid-conjugate base pairs. Critical for maintaining stable pH in biological systems and many industrial processes.
Carbon Dioxide
CO₂ dissolves in water forming carbonic acid, lowering pH. Important in ocean acidification, carbonated beverages, and respiratory regulation of blood pH. Exposed solutions slowly equilibrate with atmospheric CO₂.
pH in Biological Systems
Blood pH
Maintained at 7.35-7.45 by carbonate-bicarbonate buffer system. Kidneys and lungs regulate pH by controlling CO₂ and H⁺/HCO₃⁻ levels. pH imbalances cause serious medical conditions requiring immediate treatment.
Stomach pH
Gastric juice has pH 1-2 from hydrochloric acid secretion. Activates digestive enzymes and kills ingested bacteria. Protected by mucus layer preventing damage to stomach lining.
Cellular pH
Cytoplasm typically maintains pH 7.0-7.4. Enzyme function highly pH-dependent. Cells use membrane pumps and buffers to regulate internal pH despite metabolic acid production.
Skin pH
Healthy skin has slightly acidic pH (4.5-5.5) forming "acid mantle" protecting against pathogens. Soaps and cleansers can disrupt this pH, affecting skin health and microbial balance.
pH in Environmental Science
Acid Rain
Formed when sulfur dioxide and nitrogen oxides react with water, producing sulfuric and nitric acids. Natural rain has pH ~5.6 due to dissolved CO₂. Acid rain (pH < 5) damages ecosystems, buildings, and infrastructure.
Ocean Acidification
Increasing atmospheric CO₂ dissolves in oceans, lowering pH from historical 8.2 to current 8.1. Threatens marine organisms, especially those with calcium carbonate shells. Represents major environmental challenge.
Lake and River pH
Freshwater pH varies from 4 (acidic bog) to 9 (alkaline lake). Most aquatic life requires pH 6.5-8.5. Acid mine drainage and industrial pollution can severely alter natural water pH.
Common pH Misconceptions
Alkaline Water Health Claims
Body tightly regulates blood pH regardless of water intake. Drinking alkaline water doesn't significantly affect body pH. Stomach acid neutralizes alkaline water immediately upon consumption.
pH and Nutritional Value
Food pH doesn't determine if it's "acidifying" in the body. Lemon juice is acidic (pH 2) but metabolizes to alkaline compounds. Body pH regulation depends on metabolic processes, not dietary pH.
More Acidic = More Dangerous
Danger depends on concentration and pH. Concentrated weak acids (pH 2) can be safer than dilute strong acids (pH 4) because of different chemical properties and buffering capacity.
Tips for Accurate pH Measurement
Calibrate pH Meters Regularly
Use at least two buffer solutions (typically pH 4, 7, and 10) for calibration. Calibrate before each use for maximum accuracy. Store electrodes properly in storage solution, never dry or in distilled water.
Consider Temperature
Measure pH at consistent temperature or use temperature compensation. Note temperature when reporting pH values. Remember that pH definition itself is temperature-dependent.
Avoid Contamination
Rinse electrodes and equipment with distilled water between samples. Use clean glassware. Don't touch pH paper with fingers as skin oils affect readings.
Sample Preparation
Mix samples thoroughly before measuring. Remove particulates that can interfere with electrodes. For colored or turbid samples, consider using pH paper instead of meters.
Study Tips for Students
Understanding Logarithms
pH is logarithmic, so each unit represents 10-fold change. pH 3 has 10× more H⁺ than pH 4, and 100× more than pH 5. Practice converting between pH and [H⁺] using log/antilog functions.
Memorize Key Values
Know that pH 7 = [H⁺] = 1.0 × 10⁻⁷ M. At pH 7, [H⁺] = [OH⁻]. pH + pOH = 14. Kw = 1.0 × 10⁻¹⁴. These relationships solve most pH problems.
Practice with Real Examples
Calculate pH of household items: vinegar, baking soda solution, antacids. Measure pH of rain, tap water, juices. Connecting theory to real substances aids understanding and retention.
Understand Buffer Systems
Learn Henderson-Hasselbalch equation for buffer calculations. Understand how weak acid-conjugate base pairs resist pH changes. Critical for biochemistry and advanced chemistry courses.
Advanced pH Concepts
Activity vs Concentration
pH technically measures H⁺ activity, not concentration. In dilute solutions, activity ≈ concentration. In concentrated solutions, activity coefficients must be considered for accurate pH determination.
pH in Non-Aqueous Solvents
pH concept extends to non-aqueous systems with modified definitions. Important in organic chemistry and industrial processes using mixed or non-aqueous solvents.
pH at Different Temperatures
Kw increases with temperature: 1.14 × 10⁻¹⁵ (0°C) to 5.48 × 10⁻¹⁴ (50°C). Affects neutral pH and all pH calculations. Temperature compensation critical for accurate measurements.
Frequently Asked Questions
Why is pH measured on a 0-14 scale?
The 0-14 range reflects typical [H⁺] concentrations in aqueous solutions (1 M to 10⁻¹⁴ M). However, pH can theoretically exceed this range in very concentrated acids (negative pH) or bases (pH > 14).
Can pH be negative or greater than 14?
Yes! Concentrated acids (like 10 M HCl) have negative pH values. Concentrated bases can have pH > 14. These represent extremely corrosive solutions outside typical laboratory ranges.
Why does water have pH 7?
Pure water undergoes self-ionization: H₂O ⇌ H⁺ + OH⁻. At equilibrium (25°C), [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ M, giving pH = -log(10⁻⁷) = 7.
How accurate are pH meters?
Quality pH meters measure to ±0.01 pH units when properly calibrated. Accuracy depends on calibration, electrode condition, temperature compensation, and sample characteristics. Industrial meters can achieve ±0.001 pH units.
What's the difference between pH and acidity?
pH measures H⁺ ion concentration (intensity). Acidity (total acidity) measures total acid content that can donate H⁺ (capacity). A weak acid solution may have high total acidity but moderate pH.
Does pure water always have pH 7?
Only at 25°C. At 0°C, pure water has pH 7.5. At 50°C, pH 6.5. Temperature affects water's ion product (Kw), changing neutral pH. Always specify temperature with pH measurements.
How do buffers work?
Buffers contain weak acid and its conjugate base (or weak base and conjugate acid). Added H⁺ reacts with conjugate base; added OH⁻ reacts with weak acid. This consumes added ions, minimizing pH changes.
Why is blood pH so important?
Enzymes and proteins function optimally at specific pH. Blood pH outside 7.35-7.45 range denatures proteins, disrupts metabolic processes, and can be fatal. Body has multiple systems maintaining pH homeostasis.
Start Calculating pH Now
Use our free pH calculator to instantly compute pH, pOH, [H⁺], and [OH⁻] values. Whether you're a chemistry student learning acid-base concepts, a lab technician measuring solution pH, or monitoring pool water, our calculator provides accurate results with detailed explanations. No registration required - start calculating pH values right away!