Here are concise revision notes on Basic Chemistry, tailored for the JKSSB Forester and similar exams, focusing on 12th standard foundational concepts.
Basic Chemistry: Revision Notes
Chemistry is the study of matter and its properties, and how matter changes. It is fundamental to understanding the world around us.
I. Structure of Atom
- Atom: The smallest unit of matter that retains an element’s chemical identity.
- Composed of subatomic particles: protons, neutrons, electrons.
- Protons (p⁺): Positively charged, located in the nucleus. Mass ≈ 1 amu.
- Atomic Number (Z): Number of protons. Defines the element. (Z = p⁺)
- Neutrons (n⁰): No charge (neutral), located in the nucleus. Mass ≈ 1 amu.
- Electrons (e⁻): Negatively charged, orbit the nucleus in shells/energy levels. Mass ≈ 1/1836 amu (negligible).
- In a neutral atom, number of electrons = number of protons (e⁻ = p⁺).
- Nucleus: Dense central part of an atom containing protons and neutrons.
- Mass Number (A): Total number of protons and neutrons in the nucleus. (A = p⁺ + n⁰)
- Representation of an element: $\text{^A_Z X}$ , where X is the element symbol.
Key Highlights:
- Isotopes: Atoms of the same element (same Z) but different mass numbers (different n⁰). E.g., $\text{^1_1 H}$, $\text{^2_1 H}$ (Deuterium), $\text{^3_1 H}$ (Tritium).
- Isobars: Atoms of different elements (different Z) but same mass number (A). E.g., $\text{^40_{18} Ar}$ and $\text{^40_{20} Ca}$.
- Isotones: Atoms of different elements with the same number of neutrons. E.g., $\text{^39_{19} K}$ (n⁰=20) and $\text{^40_{20} Ca}$ (n⁰=20).
- Ions: Charged atoms or molecules.
- Cation: Positively charged (loses electrons). E.g., Na⁺.
- Anion: Negatively charged (gains electrons). E.g., Cl⁻.
Electronic Configuration:
- Distribution of electrons in various energy shells/orbitals around the nucleus.
- Bohr-Bury Rule:
- Maximum number of electrons in a shell is $\text{2n^2}$ (where n = shell number).
- K (n=1): $2(1^2)=2$
- L (n=2): $2(2^2)=8$
- M (n=3): $2(3^2)=18$
- N (n=4): $2(4^2)=32$
- Outermost shell cannot have more than 8 electrons (octet rule).
- Penultimate (second last) shell cannot have more than 18 electrons.
- Shells are filled in a stepwise manner.
- Valence Electrons: Electrons in the outermost shell. Determine chemical properties and bonding.
II. Chemical Bonding
- Chemical Bond: Force that holds atoms together in molecules or compounds.
- Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight electrons in their outermost shell (like noble gases).
- Ionic Bond (Electrovalent Bond):
- Formed by complete transfer of electrons from one atom (metal) to another (non-metal).
- Results in the formation of ions (cation and anion), which are held together by strong electrostatic forces.
- Characteristics: High melting/boiling points, soluble in polar solvents (water), conduct electricity in molten or aqueous state.
- Example: NaCl ($\text{Na⁺Cl⁻}$), MgO ($\text{Mg^{2+}O^{2-}}$).
- Covalent Bond:
- Formed by mutual sharing of electrons between two atoms (typically non-metals).
- Types:
- Single Bond: Sharing of one pair of electrons (e.g., H-H in $\text{H_2}$).
- Double Bond: Sharing of two pairs of electrons (e.g., O=O in $\text{O_2}$).
- Triple Bond: Sharing of three pairs of electrons (e.g., N≡N in $\text{N_2}$).
- Characteristics: Lower melting/boiling points than ionic compounds, often insoluble in water, do not conduct electricity.
- Example: $\text{H_2O}$, $\text{CO_2}$, $\text{CH_4}$.
- Coordinate Bond (Dative Bond):
- A type of covalent bond where both shared electrons are contributed by only one atom.
- Represented by an arrow (→) pointing from the donor atom to the acceptor atom.
- Example: Ammonium ion ($\text{NH_4⁺}$), Hydronium ion ($\text{H_3O⁺}$), formation of $\text{SO_2}$.
III. Periodic Classification of Elements
- Periodic Table: A tabular arrangement of chemical elements, ordered by their atomic number, electron configuration, and recurring chemical properties.
- Mendeleev’s Periodic Law: Properties of elements are a periodic function of their atomic masses. (Limitations led to discovery of Modern Periodic Law)
- Modern Periodic Law (Moseley): Properties of elements are a periodic function of their atomic numbers.
Key Features of Modern Periodic Table:
- Periods (Horizontal Rows): 7 periods.
- Elements in a period have the same number of electron shells.
- Properties change gradually across a period (left to right).
- Groups (Vertical Columns): 18 groups.
- Elements in a group have the same number of valence electrons (same outermost shell configuration).
- Hence, they exhibit similar chemical properties.
General Trends in the Periodic Table (Left to Right across a Period; Top to Bottom down a Group):
| Property | Across a Period (Left to Right) | Down a Group (Top to Bottom) |
|---|---|---|
| Atomic Radius | Decreases | Increases |
| Ionization Energy | Increases (more energy needed to remove e⁻) | Decreases (less energy needed to remove e⁻) |
| Electron Affinity | Increases (attracts e⁻ more) | Decreases (attracts e⁻ less) |
| Electro-negativity | Increases | Decreases |
| Metallic Character | Decreases (metals become less metallic) | Increases |
| Non-metallic Character | Increases | Decreases |
- Atomic Radius: Distance from the nucleus to the outermost electron shell.
- Ionization Energy: Energy required to remove the most loosely bound electron from an isolated gaseous atom.
- Electron Affinity: Energy released when an electron is added to an isolated gaseous atom.
- Electronegativity: Tendency of an atom to attract shared electron pairs in a covalent bond. (Pauling scale is common)
Groups to Remember:
- Group 1: Alkali Metals (Li, Na, K…) – Highly reactive, soft metals.
- Group 2: Alkaline Earth Metals (Be, Mg, Ca…) – Reactive metals.
- Group 13: Boron Family
- Group 14: Carbon Family
- Group 15: Nitrogen Family (Pnictogens)
- Group 16: Oxygen Family (Chalcogens)
- Group 17: Halogens (F, Cl, Br, I…) – Highly reactive non-metals (form salts).
- Group 18: Noble Gases (He, Ne, Ar, Kr…) – Inert, full valence shell (stable).
IV. Acids, Bases, and Salts
- Understanding these is crucial for everyday chemistry.
1. Acids:
- Arrhenius concept: Substances that produce $\text{H⁺}$ ions (or $\text{H_3O⁺}$ hydronium ions) when dissolved in water.
- Bronsted-Lowry concept: Proton ($\text{H⁺}$) donors.
- Properties:
- Sour taste.
- Turn blue litmus red.
- Corrosive nature.
- React with metals to produce hydrogen gas ($\text{H_2}$). (e.g., $\text{Zn} + \text{2HCl} \rightarrow \text{ZnCl_2} + \text{H_2}$)
- React with carbonates to produce carbon dioxide gas ($\text{CO_2}$). (e.g., $\text{CaCO_3} + \text{2HCl} \rightarrow \text{CaCl_2} + \text{H_2O} + \text{CO_2}$)
- Neutralize bases.
- Examples: HCl (Hydrochloric acid), $\text{H_2SO_4}$ (Sulphuric acid), $\text{HNO_3}$ (Nitric acid), $\text{CH_3COOH}$ (Acetic acid).
- Strong Acids: Ionize completely in water (e.g., HCl, $\text{H_2SO_4}$, $\text{HNO_3}$).
- Weak Acids: Ionize partially in water (e.g., $\text{CH_3COOH}$, Carbonic acid, Phosphoric acid).
2. Bases (or Alkalis, if soluble in water):
- Arrhenius concept: Substances that produce $\text{OH⁻}$ ions when dissolved in water.
- Bronsted-Lowry concept: Proton ($\text{H⁺}$) acceptors.
- Properties:
- Bitter taste.
- Slippery or soapy feel.
- Turn red litmus blue.
- Neutralize acids.
- Examples: NaOH (Sodium Hydroxide), KOH (Potassium Hydroxide), $\text{Ca(OH)_2}$ (Calcium Hydroxide), $\text{NH_4OH}$ (Ammonium Hydroxide).
- Strong Bases: Ionize completely in water (KOH, NaOH).
- Weak Bases: Ionize partially in water ($\text{NH_4OH}$, $\text{Mg(OH)_2}$).
3. Salts:
- Formed by the reaction of an acid and a base (neutralization reaction).
- General Reaction: Acid + Base $\rightarrow$ Salt + Water ($\text{HCl + NaOH} \rightarrow \text{NaCl + H_2O}$)
- Types:
- Neutral Salts: Formed from strong acid and strong base (e.g., NaCl). pH ≈ 7.
- Acidic Salts: Formed from strong acid and weak base (e.g., $\text{NH_4Cl}$). pH < 7.
- Basic Salts: Formed from weak acid and strong base (e.g., $\text{CH_3COONa}$). pH > 7.
- Examples: NaCl (Common salt), $\text{Na_2CO_3}$ (Washing Soda), $\text{NaHCO_3}$ (Baking Soda), $\text{CaSO_4}$ (Gypsum).
4. pH Scale:
- Measures the acidity or alkalinity of a solution.
- Scale ranges from 0 to 14.
- pH < 7: Acidic solution (lower pH = stronger acid).
- pH = 7: Neutral solution.
- pH > 7: Basic (alkaline) solution (higher pH = stronger base).
- pH Indicators: Substances that change color with changes in pH (e.g., litmus, phenolphthalein, methyl orange). Universal indicator provides a range of colors for different pH values.
V. Chemical Reactions and Equations
- Chemical Reaction: A process that leads to the transformation of one set of chemical substances to another. Bonds are broken and new bonds are formed.
- Reactants: Starting substances.
- Products: Substances formed.
- Chemical Equation: Symbolic representation of a chemical reaction using chemical formulas.
- Balanced Chemical Equation: Has an equal number of atoms of each element on both sides (reactants and products), adhering to the Law of Conservation of Mass.
Types of Chemical Reactions:
- Combination Reaction (Synthesis):
- Two or more reactants combine to form a single product.
- $\text{A + B} \rightarrow \text{AB}$
- Example: $\text{2H_2(g) + O_2(g) \rightarrow 2H_2O(l)}$ (Formation of water)
- Example: $\text{CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq)}$ (Slaking of lime)
- Decomposition Reaction:
- A single reactant breaks down into two or more simpler products.
- $\text{AB} \rightarrow \text{A + B}$
- Often requires energy (heat, light, electricity).
- Thermal Decomposition: $\text{CaCO_3(s) \xrightarrow{Heat} CaO(s) + CO_2(g)}$
- Electrolytic Decomposition: $\text{2H_2O(l) \xrightarrow{Electricity} 2H_2(g) + O_2(g)}$
- Photochemical Decomposition: $\text{2AgBr(s) \xrightarrow{Light} 2Ag(s) + Br_2(g)}$
- Displacement Reaction (Single Displacement):
- A more reactive element displaces a less reactive element from its compound.
- $\text{A + BC} \rightarrow \text{AC + B}$
- Reactivity Series is important here:
- Metals: K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Ag > Au
- A metal higher in the series can displace a metal lower in the series from its salt solution.
- Example: $\text{Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)}$
- Double Displacement Reaction:
- Exchange of ions between two reactant compounds to form two new compounds.
- $\text{AB + CD} \rightarrow \text{AD + CB}$
- Often forms a precipitate (insoluble solid) or water.
- Example (Precipitation): $\text{BaCl_2(aq) + Na_2SO_4(aq) \rightarrow BaSO_4(s) + 2NaCl(aq)}$
- Example (Neutralization): $\text{HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)}$
- Redox Reactions (Reduction-Oxidation Reactions):
- Reactions involving the transfer of electrons.
- Oxidation:
- Gain of oxygen.
- Loss of hydrogen.
- Loss of electrons (OIL – Oxidation Is Loss).
- Increase in oxidation state.
- Reduction:
- Loss of oxygen.
- Gain of hydrogen.
- Gain of electrons (RIG – Reduction Is Gain).
- Decrease in oxidation state.
- Oxidizing Agent: Substance that causes oxidation (gets reduced itself).
- Reducing Agent: Substance that causes reduction (gets oxidized itself).
- Mnemonics: OIL RIG (Oxidation Is Loss, Reduction Is Gain of electrons).
- Example: $\text{CuO(s) + H_2(g) \xrightarrow{Heat} Cu(s) + H_2O(l)}$
- CuO is reduced to Cu (loss of oxygen).
- $\text{H_2}$ is oxidized to $\text{H_2O}$ (gain of oxygen).
- CuO is the oxidizing agent.
- $\text{H_2}$ is the reducing agent.
Effects of Oxidation in Everyday Life:
- Corrosion: Slow degradation of metals due to reaction with air or water (e.g., rusting of iron: $\text{Fe_2O_3.xH_2O}$).
- Rancidity: Oxidation of fats and oils in food, leading to unpleasant smell and taste. Can be prevented by antioxidants, refrigeration, nitrogen packing.
VI. Carbon and its Compounds
- Carbon (C): A non-metal, atomic number 6. Forms a vast number of compounds due to its unique properties.
Unique Properties of Carbon:
- Catenation: Ability of carbon atoms to form strong covalent bonds with other carbon atoms, forming long chains, branched chains, and rings. This leads to complex organic molecules.
- $\text{-C-C-C-}$ (Straight chain)
- $\text{-C-C(C)-C-}$ (Branched chain)
- Cyclic Structures: Benzene ring.
- Tetravalency: Carbon has 4 valence electrons and forms 4 covalent bonds to achieve a stable octet.
Allotropes of Carbon: Different structural forms of the same element that exhibit different physical and chemical properties.
- Crystalline Allotropes:
- Diamond: Hardest natural substance, insulator, tetrahedral structure, high melting point.
- Graphite: Soft, slippery, good conductor of electricity, layers of hexagonal rings, lubricant.
- Fullerenes: Cage-like spherical structure (e.g., Buckminsterfullerene, $\text{C_{60}}$), soccer-ball shape.
- Amorphous Allotropes: Coal, charcoal, lampblack.
Hydrocarbons: Compounds made up only of carbon and hydrogen.
1. Saturated Hydrocarbons (Alkanes):
- Contain only carbon-carbon single bonds.
- General formula: $\text{C_nH_{2n+2}}$
- Suffix: -ane
- Examples: Methane ($\text{CH_4}$), Ethane ($\text{C_2H_6}$), Propane ($\text{C_3H_8}$).
2. Unsaturated Hydrocarbons:
- Contain at least one carbon-carbon double or triple bond.
- Alkenes:
- Contain at least one carbon-carbon double bond.
- General formula: $\text{C_nH_{2n}}$
- Suffix: -ene
- Examples: Ethene ($\text{C_2H_4}$), Propene ($\text{C_3H_6}$).
- Alkynes:
- Contain at least one carbon-carbon triple bond.
- General formula: $\text{C_nH_{2n-2}}$
- Suffix: -yne
- Examples: Ethyne ($\text{C_2H_2}$ – acetylene), Propyne ($\text{C_3H_4}$).
Nomenclature of Simple Organic Compounds:
- Prefix: Number of carbon atoms (Meth-1, Eth-2, Prop-3, But-4, Pent-5, Hex-6).
- Suffix: Nature of carbon-carbon bonds (-ane, -ene, -yne) or functional group.
Functional Groups: Atoms or groups of atoms that are responsible for the characteristic chemical properties of organic compounds.
| Functional Group | Formula | Suffix | Example | Common Name |
|---|---|---|---|---|
| Alcohol | -OH | -ol | $\text{CH_3OH}$ | Methanol |
| Aldehyde | -CHO | -al | $\text{CH_3CHO}$ | Ethanal (Acetaldehyde) |
| Ketone | >C=O | -one | $\text{CH_3COCH_3}$ | Propanone (Acetone) |
| Carboxylic Acid | -COOH | -oic acid | $\text{CH_3COOH}$ | Ethanoic acid (Acetic acid) |
| Ether | -O- | Alkoxyalkane | $\text{CH_3OCH_3}$ | Dimethyl ether |
| Ester | -COO- | -oate | $\text{CH_3COOCH_3}$ | Methyl ethanoate |
| Haloalkane | -X (F, Cl, Br, I) | Halo- | $\text{CH_3Cl}$ | Chloromethane |
Important Reactions of Carbon Compounds:
- Combustion: Burning of carbon compounds in oxygen to produce $\text{CO_2}$, water, heat, and light.
- $\text{CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(l) + Energy}$ (fuel reaction)
- Substitution Reaction (for alkanes): Hydrogen atoms are replaced by other atoms. (Typically with halogens in the presence of UV light).
- $\text{CH_4 + Cl_2 \xrightarrow{UV light} CH_3Cl + HCl}$
- Addition Reaction (for alkenes and alkynes): Unsaturated hydrocarbons react with other substances (like $\text{H_2}$, $\text{Cl_2}$, $\text{Br_2}$, $\text{H_2O}$) across the double or triple bond to form a saturated product.
- Hydrogenation: Addition of hydrogen in presence of a catalyst (Ni, Pd, Pt). Used for converting vegetable oils to vanaspati ghee.
- $\text{CH_2=CH_2 + H_2 \xrightarrow{Ni} CH_3-CH_3}$
- Esterification: Reaction of a carboxylic acid with an alcohol in the presence of an acid catalyst (like conc. $\text{H_2SO_4}$) to form an ester and water.
- $\text{CH_3COOH + C_2H_5OH \xrightarrow{H_2SO_4} CH_3COOC_2H_5 + H_2O}$ (Sweet-smelling compounds)
Soaps and Detergents:
- Soaps: Sodium or potassium salts of long-chain carboxylic acids.
- Form scum with hard water (calcium/magnesium salts are insoluble).
- Detergents: Sodium salts of long-chain sulphonic acids or alkyl hydrogen sulphates.
- Work well with hard water as their calcium/magnesium salts are soluble.
- Micelle Formation: Both soaps and detergents form micelles around dirt/oil particles, allowing them to be washed away with water. (Hydrophobic tail surrounds dirt, hydrophilic head faces water).
This comprehensive overview covers the essentials of Basic Chemistry (12th standard) relevant for competitive exams like JKSSB Forester. Focus on understanding the concepts, key definitions, and examples provided.
