Some Basic Concepts of Chemistry Study Notes

Some Basic Concepts of Chemistry

1.1 Introduction

• Definition of Chemistry: Chemistry is the study of matter, its physical and chemical properties, and the physical and chemical changes it undergoes under different conditions.

• Chemistry as a Central Science: The knowledge of chemistry is essential for several other fields, including:

  • Physics

  • Biological sciences

  • Applied sciences

  • Earth and space sciences

• Scope of Chemistry: Chemistry affects every aspect of life, such as:

  • The air we breathe.

  • The food we eat.

  • The fluids we drink.

  • Clothing, transportation, and fuel supplies.

• Modern Science: While an ancient science, technological advancements and sophisticated instruments have expanded chemistry into applied sciences like medicine, dentistry, engineering, agriculture, and daily home use products.

1.2 Nature of Chemistry

Chemistry is traditionally classified into five main branches:

• Organic Chemistry: The study of the properties and reactions of compounds consisting of carbon.

• Inorganic Chemistry: The study of all substances that are not organic.

• Physical Chemistry: The study of principles underlying chemistry. It deals with the properties of matter, atoms, molecules, and fundamental concepts related to electrons, energies, and dynamics. It provides the basic framework for all other branches.

• Biochemistry: (Mentioned as a branch).

• Analytical Chemistry: (Mentioned as a branch).

1.2.1 Matter

• Definition: Matter is anything that occupies space and has mass.

• Classification by Chemical Composition:

  • Pure Substances: Have a definite chemical composition and the same properties regardless of their origin.

    • Elements: Pure substances that cannot be broken down into simpler substances by ordinary chemical changes. Classified into:

      1. Metals: Have lustre (shiny), conduct heat and electricity, are ductile (drawn into wire), and malleable (hammered into sheets). Examples: $gold$, $silver$, $copper$, $iron$. Mercury is a liquid metal at room temperature.

      2. Non-metals: No lustre (exceptions: $diamond$, $iodine$), poor conductors of heat and electricity (exception: $graphite$), and are brittle. Examples: $iodine$, $nitrogen$, $carbon$.

      3. Metalloids (Semi-metals): Have properties intermediate between metals and non-metals. Examples: $arsenic$, $silicon$, and $germanium$.

    • Compounds: Pure substances that cannot be broken down into simpler substances by ordinary chemical changes. They consist of two or three elements combined in a fixed proportion.

  • Mixtures: Contain two or more substances in no fixed proportions and can be separated by physical methods.

    • Homogeneous Mixtures: Solutions where the molecules of the constituent solute and solvent are uniformly mixed throughout the bulk.

    • Heterogeneous Mixtures: Mixtures where molecules are not uniformly mixed throughout the bulk. Example: A suspension of an insoluble solid in a liquid.

1.2.3 States of Matter

• Matter exists in three interconvertible states based on temperature and pressure conditions:

  • Solids: Constituent atoms or molecules (particles) are tightly held in perfect order; they possess definite shape and volume.

  • Liquids: Particles are close to each other and can move around within the liquid. They have a definite volume but no definite shape.

  • Gases: Particles are far apart compared to solids and liquids. They have no definite shape or volume.

1.3 Properties of Matter and Their Measurement

Matter has characteristic properties classified into two categories:

• Physical Properties: Measured or observed without changing the chemical composition (e.g., colour, odour, melting point, boiling point, density).

• Chemical Properties: Properties where substances undergo a chemical change, exhibiting a change in chemical composition (e.g., coal burning in air to produce $CO_2$, or magnesium wire burning in air to form $MgO$).

1.3.1 Measurement of Properties

• Quantitative Nature: Many properties (mass, length, area, pressure, volume, time) are quantitative and expressed by a number followed by a unit (e.g., $10\,m$).

• Units: Arbitrarily decided and universally accepted standards for measurement.

• Historical Systems:

  • CGS: Centimetre (length), Gram (mass), Second (time).

  • FPS: Foot, Pound, Second.

  • MKS: Metre, Kilogram, Second.

1.3.1 SI Units (International System of Units)

Established in $1960$ as a revised metric system. It is based on the decimal system. There are seven base units (Table 1.1):

1.3.2 Specific Physical Properties

• Mass and Weight:

  • Mass: Inherent property; quantity of matter in a body. Does not vary with position. SI unit: $kg$. Lab unit: $g$ ($1\,kg = 10^3\,g$).

  • Weight: Result of mass and gravitational attraction. Varies with distance from the earth's centre.

• Length: Relative quantities include atomic radius, bond length, and wavelength. Fractional units used: nanometre ($1\,nm = 10^{-9}\,m$) and picometre ($1\,pm = 10^{-12}\,m$).

• Volume: Space occupied by a three-dimensional object. SI unit: $m^3$. Common units: $1\,L = 1\,dm^3 = 1000\,mL = 1000\,cm^3$. Lab apparatus: graduated cylinder, burette, pipette, volumetric flask.

• Density: Mass per unit volume (Characteristic property).

  • Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

  • SI Unit: $kg\,m^{-3}$. CGS Unit: $g\,cm^{-3}$ (or $g\,mL^{-1}$).

• Temperature: Measure of hotness or coldness. Three scales: Celsius ($^\circ C$), Fahrenheit ($^\circ F$), and Kelvin ($K$).

  • Freezing point of water: $0^\circ C$, $32^\circ F$, $273.15\,K$.

  • Boiling point of water: $100^\circ C$, $212^\circ F$, $373.15\,K$.

  • Conversion formulas:

    • $^\circ F = \frac{9}{5}(^\circ C) + 32$

    • $K = ^\circ C + 273.15$

  • Absolute Zero: all molecular motion stops ($-273.15^\circ C$, $0\,K$, $-459.67^\circ F$).

1.4 Laws of Chemical Combination

1.4.1 Law of Conservation of Mass

• Proposed by: Antoine Lavoisier ($1743$-$1794$).

• Observation: In combustion experiments (burning phosphorus and mercury in air), the weight gained by phosphorus equaled the weight lost by the air.

• Statement: Total mass of reactants equals the total mass of products. "Mass can neither be created nor destroyed."

1.4.2 Law of Definite Proportions

• Proposed by: Joseph Proust.

• Evidence: Experiments on natural and synthetic Cupric Carbonate. Both samples contained identical percentages by weight: $51.35\%\,copper$, $9.74\%\,carbon$, and $38.91\%\,oxygen$.

• Statement: A given compound always contains exactly the same proportion of elements by weight, irrespective of the source.

1.4.3 Law of Multiple Proportions

• Proposed by: John Dalton ($1803$).

• Statement: When two elements A and B form more than one compound, the masses of B that combine with a given mass of A are in the ratio of small whole numbers.

• Example 1: Hydrogen ($2\,g$) combines with Oxygen to form Water ($16\,g$ oxygen) or Hydrogen Peroxide ($32\,g$ oxygen). Ratio: $16:32 = 1:2$.

• Example 2: Nitrogen ($14\,g$) + Oxygen ($16\,g$) $\rightarrow$ Nitric Oxide. Nitrogen ($14\,g$) + Oxygen ($32\,g$) $\rightarrow$ Nitrogen Dioxide. Ratio: $16:32 = 1:2$.

1.4.4 Gay Lussac Law of Gaseous Volume

• Proposed by: Gay Lussac ($1808$).

• Statement: When gases combine or are produced in a reaction, they do so in a simple ratio by volume if temperature and pressure are constant.

• Example: $100\,mL$ Hydrogen + $50\,mL$ Oxygen $\rightarrow$ $100\,mL$ Water vapor. Ratio: $2:1:2$.

1.4.5 Avogadro Law

• Proposed by: Avogadro ($1811$).

• Statement: Equal volumes of all gases at the same temperature and pressure contain an equal number of molecules.

• Context: Distinguished between atoms and molecules. Explained results by assuming polyatomic (diatomic) molecules.

1.5 Dalton's Atomic Theory ($1808$)

1. Matter consists of tiny, indivisible particles called atoms.

2. All atoms of a given element have identical properties (including mass); atoms of different elements differ in mass.

3. Compounds form when atoms of different elements combine in a fixed ratio.

4. Chemical reactions involve reorganization of atoms. Atoms are neither created nor destroyed.

1.6 Atomic and Molecular Masses

1.6.1 Atomic Mass

• Scale: Relative to Carbon-12 ($C^{12}$) as the standard ($1961$ agreement).

• Definition: One atomic mass unit ($amu$) is $\frac{1}{12}$ the mass of one $C^{12}$ atom.

• **$1\,amu$in grams:**$1.66056 \times 10^{-24}\,g.\n* **Modern Symbol:** Unified mass unit (u$) or Dalton ($Da).\n* **Experimental Fact:** Mass of Oxygen atom = 26.56896 \times 10^{-24}\,g$. Atomic mass =$\frac{26.56896 \times 10^{-24}\,g}{1.66056 \times 10^{-24}\,g/u} = 16.0\,u.\n\n### 1.6.2 Average Atomic Mass\n* For elements with isotopes, it is the weighted average based on relative abundance.\n* **Calculation for Carbon:**\n * C^{12}$($98.892\%$):$12.00000\,u\n * C^{13}$($1.108\%$):$13.00335\,u\n * C^{14}$($2 \times 10^{-10}\%$):$14.00317\,u\n * Average atomic mass = (12 \times 0.98892) + (13.00335 \times 0.01108) + (14.00317 \times 2 \times 10^{-12}) = 12.011\,u.\n\n### 1.6.3 Molecular Mass\n* Sum of average atomic masses of all atoms in a molecule.\n* **Example (CO_2$):**$1(12.0\,u) + 2(16.0\,u) = 44.0\,u.\n* **Example (H_2O$):**$2(1\,u) + 16\,u = 18\,u.\n\n### 1.6.4 Formula Mass\n* Used for ionic compounds like NaCl which do not form discrete molecules. It is the sum of atomic masses in the empirical formula unit.\n* **Example (NaCl$):**$23.0\,u\,(Na) + 35.5\,u\,(Cl) = 58.5\,u.\n\n# 1.7 Mole Concept and Molar Mass\n\n* **Mole (mol$):** The amount of substance containing as many entities (atoms, molecules, ions) as there are atoms in exactly$12\,g of Carbon-12.\n* **Avogadro's Constant (N_A$):**$6.0221367 \times 10^{23}\text{ entities/mol}.\n* **Molar Mass:** The mass of one mole of a substance in grams. Numerically equal to atomic/molecular mass in u.\n * **Example:** Molar mass of Oxygen atoms = 16.0\,g\,mol^{-1}$. Molar mass of$O_2$=$32.0\,g\,mol^{-1}.\n\n# 1.8 Moles and Gases\n\n* **Molar Volume:** One mole of any gas at Standard Temperature (0^\circ C$) and Pressure ($1\,atm$) occupies$22.4\,dm^3.\n* **Formula:** n = \frac{\text{Volume of gas at STP}}{22.4\,dm^3\,mol^{-1}}.\n* **Recent IUPAC change:** Standard pressure is now 1\,bar$. Under this, molar volume is$22.71\,L\,mol^{-1}.\n\n# Questions & Discussion\n\n### Can You Tell?\n* **Classification of mixtures vs pure substances:** sea water (mixture), gasoline (mixture), skin (mixture), rusty nail (mixture/impure), textbook page (mixture), diamond (pure substance).\n* **Classify as element or compound:** mercuric oxide (compound), helium gas (element), water (compound), table salt (compound), iodine (element), mercury (element), oxygen (element), nitrogen (element).\n* **Volumes of and dihydrogen vs dioxygen:** If 10$volumes of dihydrogen react with$5$volumes of dioxygen,$10$volumes of water vapour are produced ($2:1:2 ratio).\n\n### Example Problems\n* **Urea in 5.6\,g$:** Molar mass of$NH_2CONH_2 = 60\,g\,mol^{-1}$. Moles =$0.0933\,mol$. Molecules =$0.0933 \times 6.022 \times 10^{23} = 5.618 \times 10^{22}.\n* **52\,u$of Helium vs$52\,g of Helium:**\n * 52\,u$contains$13$atoms (since$1\,atom = 4\,u).\n * 52\,g$contains$13\,moles$=$78.286 \times 10^{23} atoms.\n* **Ammonia in 67.2\,dm^3$at STP:** Moles =$\frac{67.2}{22.4} = 3.0\,mol$. Molecules =$18.066 \times 10^{23}.\n\n### Exercises Summary\n* SI unit of current: Ampere.\n* Temperature same on Celsius and Fahrenheit: -40^\circ.\n* Law illustrated by N_2 + 3H_2 \rightarrow 2NH_3$($1:3:2): Gaseous volumes.\n* Mass of oxygen used for 24\,g$Carbon making$88\,g\,CO_2$:$64\,g$($88 - 24 = 64$$).