Comprehensive Study Notes: Is Matter Around Us Pure

Scientific Definition of Purity and Nature of Matter

In everyday language, the term 'pure' is often used to describe items such as milk, ghee, butter, salt, spices, or mineral juice, implying that these substances are free from adulteration. However, from a scientific perspective, these consumables are typically mixtures comprising multiple distinct substances and are therefore not considered pure. For example, milk is a complex mixture consisting of water, fats, and proteins. A scientist defines a pure substance as a form of matter in which all constituent particles are identical in their chemical nature. This means a pure substance consists of only a single type of particles and exists as a single form of matter. In the natural world, most matter exists as mixtures of two or more pure components, such as soil, minerals, and sea water.

Introduction to Mixtures and Their Classification

Mixtures are formed by the combination of more than one kind of pure form of matter, termed a substance. A key characteristic of a substance is that it cannot be separated into other kinds of matter through any physical process. For instance, while dissolved sodium chloride ($NaCl$) can be separated from water via the physical process of evaporation, the sodium chloride itself cannot be further decomposed into its chemical components by physical means. Similarly, sugar is classified as a substance because it consists of only one type of pure matter with a uniform composition throughout. On the other hand, items like soft drinks and soil are not single substances because they contain multiple components. Regardless of the source of a substance, it will consistently exhibit the same characteristic properties. Therefore, a mixture is defined as containing more than one substance.

Characteristics and Qualitative Classification of Solutions

Mixtures are classified based on the nature of their components into two primary categories: homogeneous and heterogeneous. A homogeneous mixture, also known as a solution, possesses a uniform composition throughout its bulk. Examples include salt dissolved in water and sugar dissolved in water. Activity 2.1 demonstrates that while solutions like copper sulphate in water are homogeneous, they can have a variable composition, as evidenced by differences in color intensity depending on the amount of solute added. Heterogeneous mixtures, by contrast, contain physically distinct parts and possess non-uniform compositions. Examples include mixtures of sodium chloride and iron filings, salt and sulphur, or oil and water.

A solution is specifically a homogeneous mixture of two or more substances. While commonly perceived as liquids containing dissolved solids, liquids, or gases (like lemonade or soda water), solutions can also exist as solids (alloys) or gases (air). Homogeneity in a solution occurs at the particle level; for instance, lemonade tastes identical throughout because the particles of sugar or salt are evenly distributed. Alloys represent a special case: they are mixtures of two or more metals or a metal and a non-metal that cannot be separated by physical methods. Despite this, they are considered mixtures because they retain the properties of their constituents and have variable compositions. Brass, for example, is an alloy composed of approximately $30\%$ zinc and $70\%$ copper.

Quantitative Expression of Solution Concentration

A solution consists of two main components: the solvent and the solute. The solvent is the component that dissolves the other substance and is typically present in a larger quantity. The solute is the substance being dissolved, usually present in a smaller amount. Common examples include solid-in-liquid solutions like sugar (solute) in water (solvent), 'tincture of iodine' where iodine (solid solute) is dissolved in alcohol (liquid solvent), aerated drinks where carbon dioxide ($CO_2$ gas solute) is dissolved in water (liquid solvent), and air, which is a gas-in-gas homogeneous mixture primarily containing nitrogen ($78\%$) and oxygen ($21\%$).

The concentration of a solution is determined by the relative proportion of solute and solvent. Solutions can be described as dilute, concentrated, or saturated. A saturated solution is one in which no more solute can be dissolved at a specific temperature. The maximum amount of solute that can be dissolved at that temperature is its solubility. If the solute concentration is below this level, the solution is unsaturated. Concentration is calculated using the following formulas:

1. Mass by mass percentage of a solution = $\frac{\text{Mass of solute}}{\text{Mass of solution}} \times 100$

2. Mass by volume percentage of a solution = $\frac{\text{Mass of solute}}{\text{Volume of solution}} \times 100$

In Example 2.1, a solution with $40\,g$ of common salt in $320\,g$ of water has a mass of solution equal to $40\,g + 320\,g = 360\,g$. The mass percentage is calculated as: $\frac{40}{360} \times 100 = 11.1\%$.

Heterogeneous Systems: Suspensions and Colloidal Solutions

A suspension is a non-homogeneous system where solid particles are dispersed in a liquid without dissolving. These particles remain suspended throughout the bulk of the medium and are large enough to be visible to the naked eye. Key properties of suspensions include their instability (particles settle over time), their ability to scatter light (making the beam's path visible), and the fact that the solute can be separated by filtration. Once the particles settle, the suspension 'breaks' and no longer scatters light.

A colloidal solution, or colloid, is a mixture where particles are uniformly spread. While it may appear homogeneous due to the small size of the particles, it is actually heterogeneous. Examples include milk and ink. Colloidal particles are too small to be seen with naked eyes ($1\,nm$ to $100\,nm$) but are large enough to scatter a beam of visible light. This phenomenon is known as the Tyndall effect, named after the scientist who discovered it. Instances of the Tyndall effect include sunlight passing through a dusty room or through a dense forest canopy, where mist (water droplets in air) acts as the dispersed phase. Colloids are stable, do not settle when left undisturbed, and cannot be separated by filtration; however, they can be separated by centrifugation.

Classification and Applications of Colloidal Systems

Colloids consist of a dispersed phase (the solute-like component) and a dispersing medium (the component in which the phase is suspended). They are classified based on the state of these two components. Common types include:

1. Aerosol: Liquid or solid dispersed in gas (e.g., fog, clouds, mist, smoke, automobile exhaust).

2. Foam: Gas dispersed in liquid or solid (e.g., shaving cream, rubber, sponge, pumice).

3. Emulsion: Liquid dispersed in liquid (e.g., milk, face cream).

4. Sol: Solid dispersed in liquid (e.g., milk of magnesia, mud).

5. Gel: Liquid dispersed in solid (e.g., jelly, cheese, butter).

6. Solid Sol: Solid dispersed in solid (e.g., coloured gemstones, milky glass).

Methodologies for the Separation of Mixtures: Physical Techniques

Most natural substances are mixtures that require separation to obtain pure components. Simple physical methods like handpicking, sieving, and filtration are used for many heterogeneous mixtures. For more complex mixtures, specialized techniques are required:

1. Evaporation: Used to separate a non-volatile solute (like dye) from a volatile solvent (like water). This is how the coloured component is obtained from blue or black ink.

2. Centrifugation: This technique is based on the principle that denser particles are forced to the bottom and lighter particles stay at the top when spun rapidly. It is used in diagnostic labs for blood and urine tests, in dairies to separate butter from cream, and in washing machines to squeeze water from clothes.

3. Separating Funnel: Used to separate two immiscible liquids (like oil and water). The liquids separate into layers based on their densities. This method is also used in the iron industry to remove light slag from molten iron.

4. Sublimation: Used to separate mixtures containing a sublimable volatile component (one that changes directly from solid to gas) from a non-sublimable impurity. Examples of sublimable solids include ammonium chloride ($NH_4Cl$), camphor, naphthalene, and anthracene.

Advanced Separation Techniques: Distillation and Chromatography

Chromatography (from the Greek 'kroma' meaning 'colour') is a technique used to separate solutes that dissolve in the same solvent. It was originally used for pigments but now has wider applications. In Activity 2.7, ink components are separated as water rises on filter paper; the more soluble components rise faster. Applications include separating colours in a dye, pigments from natural colours, and drugs from blood.

Distillation is employed to separate miscible liquids that boil without decomposition and have a significant difference in their boiling points (greater than $25\,K$). For instance, acetone and water can be separated this way because acetone vaporizes and condenses first. For mixtures with boiling point differences less than $25\,K$, fractional distillation is used. This process utilizes a fractionating column filled with glass beads, which provide surfaces for repeated cooling and condensation of vapours. This is essential for separating gases from air and refining petroleum products.

Separation of Gaseous Components from Air and Water Purification

Air is a homogeneous mixture of gases that can be separated by fractional distillation. The process involves compressing and cooling air by increasing pressure and decreasing temperature to produce liquid air. This liquid air is then warmed slowly in a fractional distillation column. The gases separate at different heights according to their boiling points: Oxygen boils at $-183\,^{\circ}C$ ($20.9\%$ of air volume), Argon at $-186\,^{\circ}C$ ($0.9\%$), and Nitrogen at $-196\,^{\circ}C$ ($78.1\%$). To obtain pure oxygen, other gases must be removed. $CO_2$ is removed as dry ice during the cooling process.

Water works systems in cities use several stages for purification: storage in a reservoir, sedimentation (to allow solids to settle), loading (to sediment suspended impurities), filtration (through layers of fine sand, coarse sand, and gravel), and chlorination (to kill bacteria) before the water is supplied to homes.

Crystallization and Purification of Solids

Crystallization is a chemical technique used to purify solids by separating a pure solid in the form of its crystals from a solution. This method is superior to evaporation for two reasons: some solids decompose or char (like sugar) when heated to dryness, and some impurities might remain dissolved even after filtration, contaminating the solid upon evaporation. Practical applications include the purification of salt obtained from sea water and the separation of alum (phitkari) crystals from impure samples.

Differentiation Between Physical and Chemical Changes

Physical properties include characteristics such as colour, hardness, density, melting point, and boiling point. A physical change involves an interconversion of states (like ice melting into water) without changing the chemical composition or nature of the substance. Chemical changes, or chemical reactions, involve a change in the chemical composition, resulting in the formation of new substances. For example, while water and oil are both liquids, they differ chemically in odor and inflammability; burning is a chemical change. During the burning of a candle, both physical (melting of wax) and chemical (burning of wax) changes occur simultaneously.

Classification of Pure Substances: Elements and Compounds

Pure substances are categorized as elements or compounds. Robert Boyle first used the term 'element' in 1661, and Antoine Laurent Lavoisier established the first experimental definition: an element is a basic form of matter that cannot be broken down into simpler substances by chemical reactions. Elements are divided into:

1. Metals: Lustrous, malleable, ductile, sonorous, and good conductors of heat and electricity (e.g., gold, silver, iron). Mercury is the only metal that is liquid at room temperature.

2. Non-metals: Poor conductors, non-lustrous, and display various colours (e.g., hydrogen, oxygen, iodine, carbon).

3. Metalloids: Elements with intermediate properties (e.g., boron, silicon, germanium).

Most elements are solid at room temperature; eleven are gases, and two (mercury and bromine) are liquids. Gallium and cesium become liquid at temperatures just above room temperature ($303\,K$).

A compound is a substance composed of two or more elements chemically combined in a fixed proportion. Unlike mixtures, the properties of a compound are entirely different from its constituent elements. For instance, heating iron and sulphur produces iron sulphide ($FeS$), a compound that is no longer magnetic and reacts with dilute sulphuric acid to produce hydrogen sulphide ($H_2S$) gas (smelling of rotten eggs), whereas a simple mixture of iron and sulphur retains magnetic properties and produces hydrogen ($H_2$) gas when reacted with acid.

Comparison Between Mixtures and Compounds

The distinctions between mixtures and compounds are fundamental:

1. Composition: Mixtures have variable composition, whereas compounds have a fixed composition.

2. Formation: Mixtures are formed by simply mixing components with no new substance produced; compounds result from elements reacting to form a new substance.

3. Properties: A mixture displays the properties of its constituent substances; a compound has totally new properties distinct from its constituents.

4. Separation: Constituents of a mixture can be separated by physical methods; compounds can only be separated by chemical or electrochemical reactions.

Questions and Discussion

1. What is meant by a substance? A substance is a pure single form of matter that cannot be separated into other kinds of matter by any physical process. All constituent particles of a substance are the same in their chemical nature.

2. List the points of differences between homogeneous and heterogeneous mixtures. Homogeneous mixtures have a uniform composition throughout, and their components are not visible to the naked eye. Heterogeneous mixtures have non-uniform compositions with physically distinct parts that are often visible.

3. How are sol, solution and suspension different from each other? Significant differences exist in particle size, stability, and light scattering. A solution is a stable homogeneous mixture with particles less than $1\,nm$. A sol is a stable colloidal (heterogeneous) mixture with medium-sized particles that scatter light. A suspension is an unstable heterogeneous mixture with large visible particles that settle over time.

4. To make a saturated solution, $36\,g$ of sodium chloride is dissolved in $100\,g$ of water at $293\,K$. Find its concentration at this temperature. Mass of solute = $36\,g$. Mass of solvent = $100\,g$. Mass of solution = $36\,g + 100\,g = 136\,g$. Concentration = $\frac{36}{136} \times 100 \approx 26.47\%$.

5. Classify the following as chemical or physical changes:

• Cutting of trees: Physical

• Melting of butter in a pan: Physical

• Rusting of almirah: Chemical

• Boiling of water to form steam: Physical

• Passing of electric current through water and the water breaking down into hydrogen and oxygen gases: Chemical

• Dissolving common salt in water: Physical

• Making a fruit salad with raw fruits: Physical

• Burning of paper and wood: Chemical