Comprehensive Introduction to Chemistry: Matter, Laboratory Practices, and Basic Reactions

INTRODUCTION TO CHEMISTRY AND THE STATES OF MATTER

  • Definition of Matter: Matter is anything that has weight/mass and occupies space/volume.
  • Basics Three States of Matter:
    • Solid: Examples include soil, sand, copper metal, bucket, and ice.
    • Liquid: Examples include water, petrol, ethanol/alcohol, and mercury (the only liquid metal).
    • Gas: Examples include oxygen, nitrogen, and water vapour.
  • Characteristics of Solids:
    • Particles are closely packed.
    • Possess a definite shape and volume.
    • Particles vibrate about fixed positions.
    • Heating causes particles to vibrate more vigorously, weakening bonds, increasing spacing (expansion), and eventually causing a change of state.
  • Characteristics of Liquids:
    • Flow freely because particles slide over each other due to weak interparticle forces.
    • No definite shape; they take the shape of the container.
    • Definite volume and cannot be easily squashed.
  • Characteristics of Gases:
    • Offer the least resistance and occupy a much greater volume than the same mass of solids or liquids.
    • Particles are widely spaced apart with very weak interparticle forces.
    • No fixed volume and no fixed shape; they are restricted only by the size of the container.
    • Highly compressible as particles are far apart.
    • Particles move quickly and randomly, colliding with each other and the container walls.
  • Summary Table of State Properties:
    • Solids: Motion is vibration about fixed positions; Distance is very close; Shape is fixed; Volume is fixed.
    • Liquids: Motion includes translation, rotation, and vibration (translation is limited); Distance is very close; Shape is not fixed; Volume is fixed.
    • Gases: Motion includes translation, rotation, and vibration (particles fly rapidly and collide); Distance is very far apart; Shape is not fixed; Volume is not fixed.

METALS, NON-METALS, AND CONDUCTIVITY

  • Metals:
    • General properties: Shiny, ductile (can form wires), and malleable (can be hammered into sheets).
    • They can coil without breaking.
    • Examples: Iron, Gold, Silver, Copper. Mercury is the only known liquid metal.
  • Non-metals:
    • General properties: Dull, not ductile, not malleable, and brittle (break upon coiling).
    • Examples: Charcoal, Sulphur, plastics.
  • Conductors:
    • Materials that allow electricity to flow easily because their atoms have free/delocalized electrons.
    • Metals (molten or solid) and carbon (in the form of graphite) conduct via free electrons without chemical change.
    • Electrolytes: Molten or dissolved materials containing free-moving ions (Na+Na^{+}, Cl−Cl^{-}). Their movement constitutes an electric current.
    • List of conductors: Copper, Aluminium, Platinum, Gold, Silver, Graphite, Salt solutions (NaClNaCl), Water, People, Animals, and Trees.
  • Insulators:
    • Materials where atoms do not easily free electrons, preventing flow.
    • Examples: Glass, Porcelain, Plastic, Rubber.
  • Safety Note: Electricity takes the shortest path to the ground. Since the human body is 60%60\% water, it is a good conductor. Touching a tree touched by a fallen power line can lead to electrocution. Rubber/plastic on cords acts as an insulator to force current along metal wires.

DRUGS AND MEDICINES

  • Drug: A natural or man-made substance that alters body functioning.
  • Medicine: A drug used specifically to change abnormal body functioning back to normal.
  • Prescription: Medical instructions from a doctor/pharmacist regarding the type of medicine and the timing/duration of intake.
  • Dosage: The specific quantity of a drug required for treatment.
  • Medical Language Examples:
    • 2×42 \times 4: 2 tablets/spoonfuls taken 4 times daily (24 hours) until finished.
    • 1×21 \times 2: 1 tablet/spoonful taken 2 times daily until finished.
  • Over The Counter (OTC) Drugs: Medicines available without a prescription for mild ailments (e.g., painkillers, anti-acids, cold/flu drugs).
  • Drug Misuse/Abuse:
    • Misuse: Not following the prescribed dosage.
    • Abuse: Intentional use of a drug for purposes other than intended, often to induce a false feeling of well-being, leading to loss of mental and physical coordination.
    • Illegal Drugs: Heroin, cocaine, bhang, mandrax, morphine.
    • Legal but Abused Drugs: Miraa, alcohol, tobacco, sleeping pills.

CHEMISTRY AND THE SOCIETY

  • Definition of Chemistry: The branch of science dealing with the structure, composition, properties, and behavior of matter.
  • Science Classifications:
    • Biology: Study of living things.
    • Physical Science: Study of non-living things, split into Physics (matter in relation to energy) and Chemistry (composition of matter).
  • Role of Chemistry in Daily Life:
    • Washing/Cleaning: Interaction of water, soap, and dirt.
    • Chemicals of Life: Understanding growth, respiration, and the formation of cells using carbohydrates, proteins, and vitamins.
    • Baking: Use of baking powder and heat.
    • Medicine: Discovery and testing of drugs.
    • Cruel Oil: Fractional distillation into petrol, diesel, and kerosene.
    • Manufacturing: Production of plastics, glass, fertilizers, insecticides, soaps, and cement.
    • Diagnosis: Laboratory tests for body malfunctions.
  • Careers in Chemistry: Chemical engineering, Veterinary medicine, Medical doctor, Pharmacist, Nurse, Beautician, Chemistry teacher.

CHEMISTRY LABORATORY AND APPARATUS

  • Heating Apparatus: Bunsen burner, spirit lamp, candle, gas/kerosene stove, electric heater.
  • Measuring Apparatus:
    • Volume: Pipette (fixed small volumes), Beaker (approximate), Measuring cylinder (various volumes), Burette (accurate titration), Volumetric flask (exact volumes like 250 ml250\,ml, 1 L1\,L).
    • Mass: Beam balance (pans and scales), Electronic balance (digital screen, immediate reading in grams).
    • Temperature: Thermometer (alcohol or mercury expansion; difficult to measure solids directement).
    • Time: Stopwatch/clock (hours, minutes, seconds; features start, stop, and reset buttons).
  • Other Specific Apparatus:
    • Spatula: For scooping solids (not for accurate measurement).
    • Deflagrating spoon: For scooping and heating solids.
    • Test/Boiling Tubes: For reactions/heating.
    • Conical Flask: Narrow mouth prevents spillage; holds liquids measured by other tools.
    • Round/Flat Bottomed Flasks: Used for heating; round bottomed requires a clamp for stability.
    • Support Apparatus: Tripod stand, Wire gauze (evenly distributes heat), Clamp stand (holds apparatus firmly), Test tube holder (often has a wooden handle).
    • Funnels: Filter funnel (separation), Thistle funnel (adding liquid through a stopper), Dropping funnel (tap controls rate), Separating funnel (separates immiscible liquids).
    • Gas Jar: Used to collect and store gases.
  • Laboratory Rules & Safety:
    • Enter only with permission.
    • No unauthorized experiments.
    • Do not taste chemicals or smell them directly (waft fumes with your palm).
    • Point test tube mouths away from yourself and others while heating.
    • Report all injuries immediately.
    • Wash hands before leaving.
    • Emergency: Turn off gas, use fire extinguishers for small fires, leave through emergency doors.
  • Glassware Advantages: Transparent (visible reactions), cheap, easy to clean, unreactive.

THE BUNSEN BURNER

  • History: Invented by Robert Wilhelm Bunsen in 1854.
  • Construction:
    • Base plate: Stability.
    • Jet: Entry hole for laboratory gas (butane).
    • Collar/Sleeve: Adjustable metal controlling air entry.
    • Air hole: Formed by aligning the collar hole with the chimney hole.
    • Chimney: Tall metallic rod.
    • Tubing: Connects burner to gas tap; gas has a characteristic odor to detect leaks.
  • Lighting Procedure:
    • Close air holes via the collar.
    • Connect rubber tubing.
    • Turn on the gas.
    • Ignite at the top of the chimney quickly.
  • Types of Flames:
    • Luminous Flame: Produced when air holes are closed. Incomplete combustion occurs, creating fine unburnt carbon particles that become white-hot (giving light). It is sooty, yellow, larger, quiet, and wavy. Regions: top yellow (incomplete), unburnt gas (center), blue (sides/complete).
    • Non-luminous Flame: Produced when air holes are open. Complete combustion occurs. It is blue, hotter, non-sooty, smaller, and noisy. Regions: top colorless, blue (hottest), green (complete), unburnt gas (innermost).
  • Scientific Drawing Rules: Use 2D cross-sections; use a ruler for straight edges; use freehand for curves; show supporting stands (benches, tripods).

CLASSIFICATION OF SUBSTANCES AND MIXTURES

  • Pure Substance: Contains only one substance.
  • Impure Substance/Mixture: Combination of two or more substances that can be separated by physical means.
  • Key Definitions:
    • Solvent: Liquid in which a substance dissolves (Water is the "universal solvent").
    • Solute: Substance that dissolves.
    • Solution: Uniform mixture of solute and solvent.
    • Aqueous state: State of a solute dissolved in water.
    • Suspension/Precipitate: Mixture where an insoluble solid is suspended in a liquid (particles may settle as sediments).
  • Liquid-Liquid Mixtures:
    • Miscible: Liquids that form a uniform mixture without layers (e.g., ethanol and water).
    • Immiscible: Liquids that form layers (e.g., kerosene and water). Heavier particles settle at the bottom.
  • Alloys: Uniform mixtures of two or more metals formed by solidifying molten metals.
    • Brass: Copper and Zinc.
    • Bronze: Copper and Tin.
    • Solder: Lead and Tin (low melting point).
    • Duralumin: Aluminium, Copper, Magnesium (light and corrosion-resistant).
    • Steel: Iron, Carbon, Manganese.
    • Nichrome: Nickel and Chromium.
    • German Silver: Copper, Zinc, Nickel.

METHODS OF SEPARATING MIXTURES

  • Decantation: Pouring off liquid from settled solid sediments.
  • Filtration: Passing a mixture through a porous material (filter paper). The solid trapped is the residue; the liquid passing through is the filtrate.
  • Evaporation: Heating a solution to vaporize the solvent, leaving the solid solute (e.g., recovering salt from water).
  • Simple Distillation: Separation of solute and solvent where both are collected. Involves evaporation followed by condensation in a Liebig condenser. Used for mixtures with boiling point differences $> 40^{\circ}C$.
    • Bumping prevention: Use broken porcelain or porous pot pieces for even boiling.
  • Fractional Distillation: For miscible liquids with close boiling points (e.g., Ethanol $78^{\circ}C$ and Water $100^{\circ}C$).
    • Fractionating Column: Packed with glass beads to increase surface area for condensation of the less volatile component.
    • Industrial Applications: Refining crude oil and separating air (Nitrogen $-196^{\circ}C$, Argon $-186^{\circ}C$, Oxygen $-183^{\circ}C$).
  • Separating Funnel: For immiscible liquids. The denser liquid is run out through the tap first.
  • Sublimation and Deposition:
    • Sublimation: Solid to gas directly on heating.
    • Deposition: Gas to solid directly on cooling.
    • Substances: Iodine, NH4ClNH_{4}Cl, Camphor, FeCl3FeCl_{3}, AlCl3AlCl_{3}, Benzoic acid.
  • Chromatography: Separation of components at different rates through an absorbent material.
    • Baseline: Origin spot of the dye.
    • Solvent front: Farthest distance moved by the solvent.
    • Factors: Solubility, density, and "stickiness" of the dye.
  • Solvent Extraction: Extracting oil from crushed seeds/nuts using a volatile solvent (e.g., propanone). The solvent is later evaporated.
  • Crystallization: Forming pure crystals from a saturated solution through cooling or slow evaporation.

PHYSICAL VS. CHEMICAL CHANGES AND KINETIC THEORY

  • Physical Change: Temporary, no new substance formed, and reversible. (e.g., melting ice).
  • Chemical Change: Permanent, new substance formed, and irreversible.
  • Kinetic Theory of Matter:
    • Matter is made of particles in constant motion.
    • Melting (Solid to Liquid): Heating provides energy to weaken strong bonds (Endothermic).
    • Freezing (Liquid to Solid): Cooling removes energy allowing bonds to reform (Exothermic).
    • Boiling/Evaporation (Liquid to Gas): Energy added to break bonds completely (Endothermic). Evaporation occurs only at the surface; boiling occurs throughout the bulk.
    • Condensation (Gas to Liquid): Particles lose energy and clump together (Exothermic).
  • Heating/Cooling Curves: Graph of temperature vs. time. Plateaus indicate a change of state where temperature stays constant while bonds are broken/formed.

ATOMS, ELEMENTS, AND COMPOUNDS

  • The Atom: Smallest particle with characteristic properties.
    • Sub-atomic particles: Protons (++, in nucleus), Neutrons (neutral, in nucleus), Electrons (−-, in orbits/shells).
    • Atomic Number (Z): Number of protons (also equals number of electrons in neutral atoms).
    • Mass Number (A): protons + neutrons.
    • Neutron Number (N): A−ZA - Z.
  • Elements: Pure substances of one type of atom.
    • Natural elements: Hydrogen (H) to Uranium (U).
    • Latin-derived symbols: Sodium (Natrium - Na), Potassium (Kalium - K), Copper (Cuprum - Cu), Lead (Plumbum - Pb), Silver (Argentum - Ag), Gold (Aurum - Au), Mercury (Hydrargyrum - Hg).
  • Compounds: Chemically combined elements in fixed ratios (e.g., NaClNaCl, C6H12O6C_{6}H_{12}O_{6}). Properties differ from the constituent elements.
  • Chemical Word Equations: Reactants $\rightarrow$ Products. Arrow shows direction of change.

ACIDS, BASES, AND INDICATORS

  • Definitions:
    • Acid: Turns litmus red.
    • Base: Turns litmus blue.
    • Indicator: Substance that identifies acid, base, or neutral solutions via color change.
  • Examples of Acids:
    • Natural: Citric (citrus), Tartaric (grapes), Lactic (sour milk), Ethanoic (vinegar), Methanoic (ant stings), Carbonic (fizzy drinks).
    • Mineral: Hydrochloric acid (HClHCl), Sulphuric(VI) acid (H2SO4H_{2}SO_{4}), Nitric(V) acid (HNO3HNO_{3}).
  • Bases and Alkalis:
    • Alkalis: Bases soluble in water (NaOHNaOH, KOHKOH, NH4OHNH_{4}OH).
    • Bases (Insoluble): Magnesium oxide/hydroxide (anti-acid), Calcium oxide (cement).
  • Indicator Colors:
    • Litmus: Red in acid, Blue in base.
    • Methyl orange: Red in acid, Yellow in base, Red/Orange in neutral.
    • Phenolphthalein: Colorless in acid, Pink/Purple in base, Colorless in neutral.
  • The pH Scale: Ranges from 1 to 14.
    • 1-3: Strong acid (Red).
    • 4-6: Weak acid (Orange/Yellow).
    • 7: Neutral (Green - e.g., pure water, sodium chloride).
    • 8-11: Weak base (Blue).
    • 12-14: Strong base (Purple/Dark blue).

CHEMICAL REACTIONS OF ACIDS

  • Reaction with Metals: Acid + Reactive Metal $\rightarrow$ Salt + Hydrogen gas.
    • Mg(s)+2HCl(aq)→MgCl2(aq)+H2(g)Mg_{(s)} + 2HCl_{(aq)} \rightarrow MgCl_{2(aq)} + H_{2(g)}
    • Hydrogen Test: Extinguishes a burning splint with a "pop" sound.
    • Unreactive metals (Copper, Silver, Gold) do not react with dilute acids.
  • Reaction with Carbonates: Acid + Carbonate $\rightarrow$ Salt + Water + Carbon(IV) oxide.
    • Na2CO3(s)+2HCl(aq)→2NaCl(aq)+H2O(l)+CO2(g)Na_{2}CO_{3(s)} + 2HCl_{(aq)} \rightarrow 2NaCl_{(aq)} + H_{2}O_{(l)} + CO_{2(g)}
    • CO2CO_{2} Test: Forms a white precipitate with lime water (calcium hydroxide).
  • Neutralization: Acid + Base $\rightarrow$ Salt + Water only.
    • NaOH(aq)+HCl(aq)→NaCl(aq)+H2O(l)NaOH_{(aq)} + HCl_{(aq)} \rightarrow NaCl_{(aq)} + H_{2}O_{(l)}

AIR AND COMBUSTION

  • Composition of Air:
    • Nitrogen (78.0%78.0\%).
    • Oxygen (21.0%21.0\% - the active part supporting combustion).
    • Carbon(IV) oxide (0.03%0.03\%).
    • Noble gases (1.0%1.0\%).
  • Burning Phosphorus/Copper/Candle in Air: Experiments confirm that 21%21\% of air is used up during combustion, causing water levels to rise in a closed vessel.
  • Oxygen Preparation:
    • By decomposition of Hydrogen peroxide with Manganese(IV) oxide catalyst:
    • 2H2O2(aq)→2H2O(l)+O2(g)2H_{2}O_{2(aq)} \rightarrow 2H_{2}O_{(l)} + O_{2(g)}
    • Oxygen Test: Relights a glowing splint.
  • Burning Elements in Oxygen:
    • Metals: Form basic oxides (e.g., Magnesium $\rightarrow$ white ash MgO).
    • Non-metals: Form acidic oxides (e.g., Carbon $\rightarrow$ CO2CO_{2}, Sulphur $\rightarrow$ SO2SO_{2}).
  • Reactivity Series: Order of affinity for oxygen: K>Na>Ca>Mg>Al>C>Zn>Fe>Sn>Pb>H>Cu>Hg>Ag>Au>PtK > Na > Ca > Mg > Al > C > Zn > Fe > Sn > Pb > H > Cu > Hg > Ag > Au > Pt.
  • Redox Reactions:
    • Oxidation: Gain of oxygen.
    • Reduction: Loss of oxygen.
    • Reducing Agent: The substance that takes oxygen.
    • Oxidizing Agent: The substance that provides oxygen.

WATER AND HYDROGEN

  • Water Cycle: Solar energy causes evaporation/transpiration; cooling leads to condensation (clouds) and freezing; precipitation returns as rain/snow.
  • Tests for Water:
    • Anhydrous copper(II) sulphate(VI): White $\rightarrow$ Blue.
    • Anhydrous cobalt(II) chloride paper: Blue $\rightarrow$ Pink.
    • Purity: Boiling point of $100^{\circ}C$ and Melting point of $0^{\circ}C$ at 1 atm.
  • Metal Reactions with Water:
    • Sodium: Floats, melts to a ball, reacts vigorously to form NaOHNaOH and H2H_{2}.
    • Potassium: Explodes into flames.
    • Magnesium: Reacts slowly with cold water but burns vigorously in steam to form Magnesium oxide and Hydrogen.
    • Iron: Reacts with steam to form Tri-iron tetra-oxide (Fe3O4Fe_{3}O_{4}).
  • Hydrogen Preparation: Zinc + dilute acid (HClHCl or H2SO4H_{2}SO_{4}).
    • Zn(s)+2HCl(aq)→ZnCl2(aq)+H2(g)Zn_{(s)} + 2HCl_{(aq)} \rightarrow ZnCl_{2(aq)} + H_{2(g)}
    • Uses: Hydrogenation of oils (margarine), rocket fuel, Haber process (Ammonia), weather balloons.