CHEMISTRY
PROJECT
THE pH SCALE AND ITS APPLICATIONS
Submitted by:
Name: JEYSAN MICHAEL
Class: XI A
Acknowledgements
I express my heartfelt gratitude to my
Chemistry teacher Mrs. Keerthy for her
valuable guidance in the successful
completion of this project on
‘pH Scale and Its Applications’.
I also thank my parents for their support
and encouragement throughout this work
and also grateful to my friends and
classmates who shared ideas and
resources that helped enrich this project.
Introduction to Chemistry
Chemistry is the scientific study of matter—its structure, +
properties, composition, and the transformations it undergoes.
From the smallest atomic particles to complex molecules,
chemistry explains the interactions that shape our world. Among
its most significant branches is acid-base chemistry.
Acids and bases are omnipresent: from the sour taste of citrus
fruits to the bitter feel of soap, their presence can be sensed
everywhere. Acids are substances that release hydrogen ions (H⁺)
in water, while bases release hydroxide ions (OH⁻).
The interaction between acids and bases governs vital processes in
biology, medicine, agriculture, and environmental science.
Understanding their strengths, reactions, and measurements is
crucial to both academic learning and real-life applications. The
central concept tying these together is the pH scale.
History of pH Scale
The pH scale was introduced in 1909 by Danish chemist Søren
Sørensen while working at the Carlsberg Laboratory in
Copenhagen. He needed a simple way to represent hydrogen ion
concentrations in solutions, as earlier methods involved
cumbersome calculations.
Sørensen proposed using a logarithmic scale:
pH = –log [H⁺].
This innovation allowed scientists to easily communicate whether
a solution was acidic, neutral, or basic without writing long
exponents. The logarithmic nature of the scale meant that a one-
unit change in pH represented a tenfold change in hydrogen ion
concentration. Today, more than a century later, the pH scale
remains one of the most important measurement systems in
science, bridging laboratory experiments and real-world practices
in biology, industry, and environment.
Definition of pH
The pH of a solution is defined as the negative logarithm of the
hydrogen ion concentration:
pH = –log [H⁺].
Pure water has a hydrogen ion concentration of 1 × 10⁻⁷ mol/L,
giving it a neutral pH of 7.
Values less than 7 indicate acidity, while values greater than 7
indicate alkalinity. The elegance of the pH definition lies in its
simplicity, yet its significance is profound.
It explains why lemon juice (pH ≈ 2) tastes sour, why baking soda
solution (pH ≈ 9) feels slippery, and why maintaining a stable
blood pH (≈ 7.4) is essential for human survival.
Measuring pH gives us insight into the chemical nature of a
solution, helping scientists and industries make decisions ranging
from safe drinking water standards to drug formulation and soil
treatment.
The pH Scale
The pH scale extends from 0 to 14. Strong acids such as
hydrochloric acid (HCl) have pH values close to 0–1, while strong
bases like sodium hydroxide (NaOH) show pH values near 13–14.
Weak acids, such as acetic acid in vinegar, lie around pH 4–6,
while weak bases like ammonium hydroxide occupy the range of
8–10.
A solution with pH 7 is neutral, corresponding to pure water. The
scale is not merely theoretical; it categorises substances in practical
contexts. For example, gastric juice in our stomach has pH ≈ 1.5–
2, aiding digestion, while household bleach has pH ≈ 12, giving it
strong cleaning power. This categorisation makes the scale an
indispensable tool for scientists, healthcare workers, farmers, and
engineers alike.
Strong Acids and Bases
A strong acid ionises completely in water, releasing a high
concentration of H⁺ ions. Examples include hydrochloric acid
(HCl), nitric acid (HNO₃), and sulfuric acid (H₂SO₄).
These acids have low pH values (0–3) and are used in industries
for fertilisers, cleaning agents, and laboratory reagents.
Strong bases such as sodium hydroxide (NaOH) and potassium
hydroxide (KOH) dissociate completely to release OH⁻ ions,
showing pH values between 11–14.
They find applications in soap making, paper production, and
chemical synthesis. Due to their corrosive nature, both strong acids
and bases must be handled with extreme caution. Their study helps
students appreciate the balance required in chemical usage—
powerful, yet potentially hazardous.
Weak Acids and Bases
Unlike strong acids, weak acids ionise only partially in aqueous
solutions. A classic example is acetic acid (CH₃COOH), found in
vinegar. It establishes an equilibrium between undissociated
molecules and ions, resulting in a pH between 3 and 6 depending
on concentration.
Weak bases such as ammonium hydroxide (NH₄OH) also
dissociate incompletely, producing fewer hydroxide ions compared
to strong bases. Their pH values generally fall between 8 and 10.
Though weaker in strength, these compounds are no less
important.
Weak acids contribute to food flavour and preservation, while
weak bases are used in textiles and as mild cleaning agents. Their
moderate reactivity makes them safer for everyday applications.
Electrolytes and Ionisation
Electrolytes are compounds that conduct electricity when dissolved
in water due to the formation of ions. Strong electrolytes, like HCl
or NaOH, dissociate almost completely, providing high
conductivity.
Weak electrolytes, like acetic acid or ammonia, ionise partially,
producing lower conductivity. The degree of ionisation (α) is the
ratio of ionised molecules to the total molecules present. It
influences both the conductivity and the pH of the solution.
Understanding electrolytes is crucial not only in chemistry but also
in physiology. For instance, electrolytes in human blood—sodium,
potassium, calcium—regulate muscle function and nerve
transmission, demonstrating the importance of ionisation beyond
the laboratory.
Relationship between pH, pOH, and Kw
The ionisation of water establishes a constant known as Kw, where
Kw = [H⁺][OH⁻] = 1 × 10⁻¹⁴ at 25°C.
From this relationship, we derive that
pH + pOH = 14.
This equation links the acidity and alkalinity of a solution in a
simple form. For example, if a solution has pH 4 (acidic), then
pOH = 10.
Conversely, if pH = 9 (basic), then pOH = 5.
This interconnection allows chemists to calculate hydroxide ion
concentrations from hydrogen ion concentrations and vice versa. It
is particularly useful in titrations and buffer preparations, where
maintaining precise pH is essential.
Buffer Solutions
Buffers are solutions that resist changes in pH when small amounts
of acid or base are added. They are typically composed of a weak
acid and its conjugate base or a weak base and its conjugate acid.
For example, a mixture of acetic acid and sodium acetate forms an
effective buffer around pH 4.7.
The mathematical treatment of buffers is expressed by the
Henderson–Hasselbalch equation: pH = pKa + log ([A⁻]/[HA]).
This formula allows prediction of buffer capacity and efficiency.
Biological systems heavily depend on buffers. Human blood uses
the carbonic acid–bicarbonate buffer system to maintain pH around
7.4, ensuring survival. Without buffers, even minor metabolic
changes would make life unsustainable.
Indicators
Indicators are substances that change colour depending on the pH
of the solution. Litmus, one of the oldest indicators, turns red in
acidic conditions and blue in alkaline conditions.
Methyl orange shows red at pH < 3.1 and yellow above pH 4.4,
making it useful in strong acid titrations. Phenolphthalein,
colourless below pH 8.3 and pink above pH 10, is widely used in
base titrations.
Universal indicators combine several indicators to provide a full
colour chart across pH 1 to 14. Indicators are invaluable in
laboratories and industries for quick pH detection, ensuring safety
and efficiency in chemical processes.
Applications in Biology
The pH scale is integral to biological systems. Human blood must
maintain pH around 7.35–7.45. Even slight deviations can be fatal,
as enzyme activity is highly sensitive to pH changes.
Digestive fluids also illustrate the importance of pH: gastric juice
has pH ~2, aiding protein breakdown, while pancreatic juice with
pH ~8 neutralises acidity in the small intestine. Plant growth also
depends on soil pH.
Certain nutrients are only available in specific pH ranges, directly
affecting agricultural productivity. pH also influences microbial
growth, as bacteria and fungi thrive in specific pH environments.
Thus, understanding pH is critical not only for scientists but also
for doctors, farmers, and environmentalists.
Applications in Agriculture
Soil pH determines nutrient availability for crops. Acidic soils (pH
< 6) often lack calcium and magnesium, while alkaline soils (pH >
8) may lock up iron and phosphorus.
Farmers adjust soil pH using lime (CaCO₃) to raise pH or gypsum
to lower it. Crop yields are strongly correlated with optimal pH
levels. For example, rice grows best in slightly acidic soils, while
wheat prefers near-neutral conditions.
Fertiliser effectiveness is also pH-dependent; nitrogen fertilisers
may acidify the soil over time, requiring corrective measures.
Hence, measuring and adjusting soil pH is a vital practice in
sustainable agriculture, directly influencing food security.
Applications in Industry
Industries rely on pH control for efficiency and product quality. In
pharmaceuticals, drug stability and absorption depend on pH. The
food industry uses pH adjustments in cheese making, fermentation,
and preservation.
Textile and dye industries rely on pH-sensitive dyes for uniform
colouring. Wastewater treatment plants constantly monitor pH to
neutralise harmful effluents before release. Paper manufacturing
requires strong bases, while soap and detergent industries adjust
alkalinity to control product quality.
These diverse applications show that pH control is not an abstract
concept but a cornerstone of global industrial operations.
Applications in Environment
Environmental chemistry heavily involves pH. Acid rain, with pH
below 5.6, results from sulphur dioxide and nitrogen oxides
dissolving in water, damaging crops, soil, and aquatic life.
Monitoring natural water bodies for pH is essential to assess
pollution levels. Ocean acidification, caused by CO₂ dissolving
into seawater, lowers pH and affects marine organisms like corals
and shellfish.
Forest health is also linked to soil pH, as acidic deposition leaches
nutrients. Thus, environmental conservation strategies must
include continuous monitoring and regulation of pH levels to
preserve ecosystems and biodiversity.
Experimental Work
Experiments bring theory to life. Simple pH testing of household
substances (lemon juice, milk, soap, baking soda solution) with
litmus or pH paper demonstrates acidity and alkalinity. A pH meter
provides more accuracy, showing slight variations in common
liquids.
Buffer preparation experiments, such as mixing acetic acid with
sodium acetate, show resistance to pH changes. Acid-base
titrations further illustrate pH shifts, with indicators marking
endpoints.
Such hands-on activities reinforce theoretical learning and reveal
the practical significance of the pH scale in daily life.
Solved Problems
1: Find pH of 0.01 M HCl.
Solution:
[H⁺] = 0.01 → pH = –log (0.01) = 2.
2: Calculate pH of 0.1 M CH₃COOH (Ka = 1.8 × 10⁻⁵).
Solution:
[H⁺] = √(Ka × C) = √(1.8 × 10⁻⁵ × 0.1) = 1.34 × 10⁻³; pH = 2.87.
3: Buffer solution of 0.1 M acetic acid and 0.1 M sodium acetate.
Solution:
pH = pKa + log([A⁻]/[HA]) = 4.74 + log(1) = 4.74.
Conclusion
The pH scale unifies the concepts of acidity and alkalinity, offering
a simple yet powerful tool to measure chemical reactivity. Its
applications span biology, agriculture, industry, and environment.
By mastering this concept, students appreciate chemistry’s
relevance to life. This project has attempted to present pH in depth
with both theoretical and experimental perspectives.
References:
1. ISC Chemistry Syllabus
2. Byju’s Learning App
3. Jagran Josh Education
4. Carlsberg Laboratory publications
