SSLC:Chemistry Notes
Chemistry: Standard-X Study Notes
Unit 1: Nomenclature of Organic Compounds and Isomerism
This unit explores the vast world of organic compounds, focusing on their structure, naming (nomenclature), and isomerism. Organic chemistry is a specific branch of chemistry dealing with billions of carbon compounds that combine with elements like hydrogen, oxygen, nitrogen, and others, proving indispensable in modern life.
Nomenclature of Alkanes with One Branch Branched hydrocarbons follow specific IUPAC rules for naming:
• Main Chain Selection: The longest chain containing the maximum number of carbon atoms is considered the main chain, with the remaining parts as branches.
• Numbering Rule: Numbering should be done so that the carbon atom carrying the branch receives the lowest possible number. This can be done by numbering from either the left or the right of the chain.
• Alkyl Groups: Small branches attached to carbon atoms are called alkyl groups. An alkyl group is formed when a hydrogen atom is removed from a saturated hydrocarbon.
◦ To name an alkyl group, 'yl' is added to the word root of the corresponding alkane.
◦ Examples: Methyl (–CH3), Ethyl (–CH2–CH3), Propyl (–CH2–CH2–CH3).
• IUPAC Naming Format: The IUPAC name follows the format: Position number of branch + hyphen + name of alkyl group + word root + suffix (ane). A hyphen (–) separates numerals and alphabets.
◦ Example: For CH3 – CH2 – CH2 – CH – CH3 with a CH3 branch at position 2, the IUPAC name is 2–Methylpentane.
Nomenclature of Alkanes with More Than One Branch
• Longest Chain and Numbering: Select the longest carbon chain. Number the carbon atoms from left to right or right to left to ensure the carbon atoms with branches get the lowest position numbers.
◦ Example: For CH3 – CH – CH2 – CH – CH2 – CH3 with methyl branches at positions 2 and 4 (from left to right numbering), the correct position numbers are 2, 4.
• Prefixes for Multiple Branches: If the same branch appears more than once, use prefixes like di- (two), tri- (three), tetra- (four) to indicate the number of branches. Position numbers should be separated by commas.
◦ Example: For a 7-carbon chain with methyl branches at positions 2 and 5 (right to left numbering), the IUPAC name is 2,5–Dimethylheptane.
• Tie-breaking Rule for Numbering: If the carbon atom containing the first branch gets the same position number when numbered from either side, numbering should be done so that the carbon atom containing the second branch gets a lower position number.
◦ Example: For a hexane chain with methyl branches at 2,3,5, the numbering that gives 2,3,5 is preferred over 2,4,5. The IUPAC name is 2,3,5–Trimethylhexane.
• Repeating Position Numbers: If a carbon atom has two identical branches, their position numbers should be repeated.
◦ Example: For a propane chain with two methyl branches at carbon 2, the IUPAC name is 2,2–Dimethylpropane.
Writing Structural Formulae from IUPAC Names To write the structural formula from an IUPAC name (e.g., 2,3-Dimethylpentane):
1. Identify the main chain (pentane = 5 carbon atoms).
2. Draw the carbon backbone.
3. Identify the branches (Dimethyl = two methyl groups).
4. Identify their position numbers (2,3).
5. Attach the branches to the main chain at the specified positions.
6. Complete the structure by filling all carbon valencies with hydrogen atoms.
Nomenclature of Unsaturated Hydrocarbons These include alkenes (with double bonds) and alkynes (with triple bonds).
• Numbering Rule: The numbering should be done so that the carbon atoms linked by the double or triple bond receive the lowest position number.
• Alkenes Naming Format: Word root + hyphen + position of double bond + hyphen + suffix (ene).
◦ Example: CH2 = CH – CH2 – CH3 (4 carbons, double bond at position 1) is But–1–ene.
◦ Example: CH3 – CH = CH – CH3 (4 carbons, double bond at position 2) is But–2–ene.
• Alkynes Naming Format: Word root + hyphen + position of triple bond + hyphen + suffix (yne).
◦ Example: CH ≡ C – CH2 – CH3 (4 carbons, triple bond at position 1) is But–1–yne.
◦ Example: CH3 – C ≡ C – CH3 (4 carbons, triple bond at position 2) is But–2–yne.
• Double and Triple Bonds as Functional Groups: The position of double and triple bonds are crucial for naming and are considered functional groups.
Functional Groups An atom or a group of atoms bonded to carbon in an organic compound that determines its distinctive chemical and physical properties is called a functional group.
1. Hydroxyl Group (–OH)
◦ Aliphatic hydrocarbons with an –OH group are called alcohols.
◦ Naming: Replace 'e' in the corresponding alkane name with 'ol'.
▪ Format: Alkane – e + ol → Alkanol.
▪ Example: Methanol (CH3–OH), Ethanol (CH3–CH2–OH).
◦ Position Specification: When there are more than two carbon atoms, the position of the –OH group must be specified.
▪ Main chain contains the –OH group.
▪ Carbon atoms are numbered so the functional group gets the lowest position number.
▪ Format: Alkane – e + hyphen + position number of –OH group + hyphen + ol.
▪ Example: CH3 – CH2 – CH2 – OH is Propan–1–ol.
▪ Example: CH3 – CH(OH) – CH3 is Propan–2–ol.
2. Carboxyl Group (–COOH)
◦ Compounds containing this group are carboxylic acids.
◦ Naming: Replace the last letter 'e' of the corresponding alkane with 'oic acid'. The carbon atom in the carboxyl group is included in the main chain count.
▪ Format: Alkane – e + oic acid → Alkanoic acid.
▪ Example: H–COOH is Methanoic acid.
▪ Example: CH3–COOH is Ethanoic acid.
◦ Fatty Acids: Saturated or unsaturated carboxylic acids with long aliphatic chains (e.g., palmitic acid, stearic acid). They are used in soap, detergent, and cosmetic manufacturing.
3. Aldehyde Group (–CHO)
◦ Compounds with this group are called aldehydes.
◦ Naming: Replace the last letter 'e' in the corresponding alkane name with 'al'. The carbon atom in the aldehyde group is included in the main chain count.
▪ Format: Alkane – e + al → Alkanal.
▪ Example: CH3–CHO is Ethanal.
4. Keto Group (C=O)
◦ Compounds with this group are called ketones.
◦ Naming: Replace the letter 'e' of the corresponding alkane with 'one'. The carbon atom in the keto group is included in the main chain count.
▪ Format: Alkane – e + one → Alkanone.
▪ Example: H3C–CO–CH3 is Propanone (also known as acetone).
◦ Position Specification: The position of the functional group must be considered for ketones with more than 3 carbon atoms.
▪ Example: CH3–CH2–CH2–CO–CH3 is Pentan–2–one.
5. Halo Group (–F, –Cl, –Br, –I)
◦ Compounds formed by replacing hydrogen atoms in hydrocarbons with halogen atoms are halo compounds.
◦ Nomenclature: Position of the halo group + hyphen + name of the halo group + name of the alkane.
▪ Example: CH3–CH2–CH2–Cl is 1–Chloropropane.
6. Alkoxy Group (–O–R)
◦ Compounds containing the alkoxy group are called ethers. 'R' denotes an alkyl group.
◦ Nomenclature: Ethers are named as alkoxyalkane. The longer alkyl group on either side of the ether linkage (–O–) is considered the alkane, and the shorter is the alkoxy group.
▪ Example: CH3–O–CH2–CH3 is Methoxyethane.
▪ Example: CH3–O–CH3 is Methoxymethane.
Aromatic Compounds
• These are cyclic compounds like benzene (C6H6).
• Derivatives like Phenol (C6H5–OH, benzene with an –OH group) and Benzoic acid (C6H5–COOH, benzene with a –COOH group) are important.
Isomerism Isomers are compounds having the same molecular formula but different chemical and physical properties due to differing structural formulae. This phenomenon is called isomerism.
1. Chain Isomerism
◦ Compounds with the same molecular formula but different structures of the carbon chain.
◦ Example: Butane (CH3–CH2–CH2–CH3) and 2-Methylpropane (CH3–CH(CH3)–CH3) are chain isomers.
2. Position Isomerism
◦ Compounds having the same molecular formula and the same functional group, but differing in the position of the functional group. This also applies to the position of double or triple bonds.
◦ Example: Propan–1–ol (CH3–CH2–CH2–OH) and Propan–2–ol (CH3–CH(OH)–CH3) are position isomers.
◦ Example: But–1–ene (CH3–CH2–CH=CH2) and But–2–ene (CH3–CH=CH–CH3) are position isomers.
3. Functional Isomerism
◦ Compounds having the same molecular formula but different functional groups.
◦ Example: Ethanol (CH3–CH2–OH, alcohol) and Methoxymethane (CH3–O–CH3, ether) are functional isomers.
◦ Example: Propanal (CH3–CH2–CHO, aldehyde) and Propanone (CH3–CO–CH3, ketone) are functional isomers.
4. Metamerism
◦ A type of isomerism where compounds have the same molecular formula but different alkyl groups attached to either side of a bivalent functional group (a group having a valency of 2, e.g., –O–, C=O).
◦ Example: Diethyl ether (CH3–CH2–O–CH2–CH3) and Methyl propyl ether (CH3–O–CH2–CH2–CH3) are metamers.
◦ Example: Pentan–3–one (CH3–CH2–CO–CH2–CH3) and Pentan–2–one (CH3–CO–CH2–CH2–CH3) are also metamers.
--------------------------------------------------------------------------------
Unit 2: Chemical Reactions of Organic Compounds
Organic reactions are chemical reactions involving carbon compounds. Major types include substitution, addition, combustion, thermal cracking, and polymerisation.
1. Substitution Reactions
• Definition: Chemical reactions where an atom or group of atoms in a compound is replaced with another atom or group of atoms.
• Example: The reaction of methane (CH4) with chlorine (Cl2) in the presence of sunlight.
◦ CH4 + Cl2 → CH3Cl + HCl (Chloromethane is formed).
◦ Further reactions can occur: CH3Cl + Cl2 → CH2Cl2 + HCl, then CHCl3, and finally CCl4.
2. Addition Reactions
• Definition: Reactions in which unsaturated organic compounds (with double or triple bonds) combine with certain molecules to form saturated compounds. Reactions where triple-bonded compounds partially combine to form double-bonded compounds are also addition reactions.
• Characteristic: The carbon-carbon double or triple bond changes.
• Examples:
◦ Ethene (C2H4) + Hydrogen (H2) → Ethane (C2H6) (in presence of Nickel catalyst at high temperature).
◦ Ethyne (H–C≡C–H) + H2 → Ethene (CH2=CH2).
◦ Ethene (CH2=CH2) + Cl2 → 1,2–Dichloroethane (CH2Cl–CH2Cl).
3. Polymerisation
• Definition: The process by which simple molecules (monomers) join together to form large, complex molecules (polymers).
• Types:
◦ Addition Polymers: Obtained by the repeated addition reaction of monomers.
▪ Example: Ethene (monomer) → Polyethene (polythene, polymer).
▪ Polyvinyl Chloride (PVC) is formed from vinyl chloride.
▪ Polytetrafluoroethene (Teflon) is formed from tetrafluoroethene. It is used for non-stick cookware due to its high temperature resistance.
▪ Other examples: Natural rubber (from isoprene), Orlon (from acrylonitrile).
◦ Condensation Polymers: Formed when different monomers combine, accompanied by the removal of simple molecules (e.g., H2O).
▪ Example: Nylon 66 is obtained from the condensation polymerisation of adipic acid and hexamethylenediamine at high temperature and pressure, with the removal of H2O.
▪ Other examples: Phenol formaldehyde resin (bakelite) from phenol and formaldehyde, Polyethylene terephthalate (polyester) from ethylene glycol and terephthalic acid.
4. Thermal Cracking
• Definition: When heated in the absence of air, some hydrocarbons with high molecular weight decompose into hydrocarbons with lower molecular weight.
• Products: Can include both saturated and unsaturated hydrocarbons, depending on temperature, pressure, and the nature of the hydrocarbon.
• Environmental Benefit: Can break down plastic wastes into lighter molecules, helping to control pollution.
• Example: Propane (CH3–CH2–CH3) → Methane (CH4) + Ethene (CH2=CH2) when heated.
5. Combustion of Hydrocarbons
• Definition: When hydrocarbons burn, they combine with oxygen in the air to form carbon dioxide (CO2) and water (H2O), along with heat and light.
• Products: All hydrocarbons produce CO2 and H2O upon complete combustion.
• Examples:
◦ Methane (CH4) + Oxygen (O2) → CO2 + H2O + Heat.
◦ Butane (C4H10) + O2 → CO2 + H2O + Heat.
• Environmental Impact: Burning fossil fuels (like methane, butane) releases CO2, highlighting the need to control their use.
Some Important Organic Compounds
Methanol (CH3–OH)
• Also called wood spirit.
• Industrial Production: Treating carbon monoxide (CO) with hydrogen (H2) in the presence of catalysts at 573 K.
• Properties: Poisonous substance.
• Uses: Manufacturing varnish, paint, formic acid, formaldehyde (40% solution is formalin).
Ethanol (CH3–CH2–OH)
• Also called ethyl alcohol.
• Industrial Preparation: Fermentation of molasses (viscous sugar solution from sugarcane production).
◦ Yeast is added to dilute molasses.
◦ Invertase enzyme (from yeast) converts sugar solution to glucose and fructose.
◦ Zymase enzyme (from yeast) converts glucose and fructose into ethanol and CO2.
◦ C12H22O11 + H2O (with Invertase) → C6H12O6 (Glucose) + C6H12O6 (Fructose).
◦ C6H12O6 (Glucose/Fructose) (with Zymase) → 2C2H5OH (Ethanol) + 2CO2.
• Forms of Ethanol:
◦ Wash: 8-10% ethanol obtained from fermentation.
◦ Rectified spirit: 95.6% ethanol obtained by fractional distillation of wash.
◦ Absolute alcohol: 100% ethanol.
◦ Power alcohol: Mixture of 20% absolute alcohol and 80% petrol, used as vehicle fuel.
◦ Denatured spirit: Ethanol made toxic by adding substances like methanol, pyridine, or rubber distillate to prevent misuse as a beverage.
◦ Methylated spirit: Denatured spirit specifically treated with methanol.
• Uses: Production of power alcohol, solvent for medicines, manufacturing paints, preservatives, production of organic compounds.
Ethanoic Acid (CH3–COOH)
• Also called acetic acid.
• Industrial Preparation:
◦ Treating methanol (CH3–OH) with carbon monoxide (CO) in the presence of a catalyst.
◦ Fermentation of ethanol with acetobacter bacteria in the presence of air yields 5-8% ethanoic acid, known as vinegar.
• Uses: Manufacturing vinegar, production of acetic anhydride, acetate ester, synthetic fibres, solvent for polymers and resins, disinfectants, medicines.
Esters
• Formation: Formed when alcohols react with carboxylic acids, a reaction called esterification.
• General Formula: R – C(=O) – O – R1 (where R, R1 are alkyl groups).
• Example: Ethanoic acid (CH3–COOH) + Ethanol (CH3–CH2–OH) → Ethyl ethanoate (CH3–COO–CH2–CH3) + H2O (in presence of concentrated H2SO4).
• Properties & Uses: Esters have the fragrance of flowers and fruits and are used to make artificial perfumes and juices.
◦ Example: Isoamyl acetate (banana fragrance), Benzyl ethanoate (jasmine fragrance), Octyl ethanoate (orange fragrance), Ethyl butanoate (pineapple fragrance).
• Methyl Salicylate: A methyl ester of salicylic acid, used to relieve joint and muscle pain, also known as oil of wintergreen due to its extraction from evergreen plants.
Medicines Chemistry plays a pivotal role in medical science and pharmaceutical research.
• Analgesics: Relieve pain (e.g., Aspirin, Paracetamol).
• Antipyretics: Reduce body temperature (e.g., Paracetamol).
• Antiseptics: Control microorganisms (e.g., Dettol).
• Antibiotics: Destroy infectious microorganisms and prevent their growth (e.g., Penicillin).
• Paracetamol: N-acetyl-p-amino phenol, commonly used as an antipyretic and analgesic with fewer side effects, though high consumption can affect the liver. Included in WHO essential medicines list.
• Aspirin: Chemical name is acetyl salicylic acid. Used as an analgesic and has anti-blood coagulant properties to prevent heart attack.
--------------------------------------------------------------------------------
Unit 3: Periodic Table and Electron Configuration
This unit delves into the arrangement of elements based on their chemical properties and the distribution of electrons within atoms.
Limitations of Bohr Model and Quantum Mechanical Model
• Bohr model proposed electrons in definite energy orbits.
• Later studies (Louis de Broglie's wave nature of matter, Heisenberg's Uncertainty Principle) showed that electrons cannot be considered merely as particles moving in orbits, as it's impossible to simultaneously determine their exact position and velocity.
• The quantum mechanical model was developed, stating that electrons are found in regions around the nucleus where there is a maximum probability of finding them, called orbitals.
• Quantum numbers are used to describe the characteristics of orbitals and electrons.
Quantum Numbers
1. Principal Quantum Number (n):
◦ Represents shells or principal energy levels.
◦ Values: n = 1, 2, 3, 4, ....
◦ Corresponds to shells: n=1 (K shell), n=2 (L shell), n=3 (M shell), n=4 (N shell).
2. Azimuthal Quantum Number (l):
◦ Defines the three-dimensional shape of the orbital.
◦ Represents subshells (s, p, d, f).
◦ Values: Ranges from 0 to (n-1).
▪ l = 0 denotes s subshell (sharp, spherical shape, 1 orbital).
▪ l = 1 denotes p subshell (principal, dumbbell shape, 3 orbitals).
▪ l = 2 denotes d subshell (diffuse, 5 orbitals).
▪ l = 3 denotes f subshell (fundamental, 7 orbitals).
◦ The number of subshells in a particular shell equals 'n'.
3. Magnetic Quantum Number (m):
◦ Represents the difference in the orientation of orbitals.
◦ Values: For a particular 'l', there are (2l + 1) values for 'm'.
▪ For l=0 (s subshell), m has 1 value (1 orbital).
▪ For l=1 (p subshell), m has 3 values (3 orbitals).
▪ For l=2 (d subshell), m has 5 values (5 orbitals).
▪ For l=3 (f subshell), m has 7 values (7 orbitals).
Number of Orbitals and Electrons
• Total number of orbitals in each shell = n².
◦ K shell (n=1): 1² = 1 orbital.
◦ L shell (n=2): 2² = 4 orbitals.
• Maximum number of electrons that can be accommodated in each shell = 2n².
◦ K shell (n=1): 2(1)² = 2 electrons.
◦ L shell (n=2): 2(2)² = 8 electrons.
• Maximum number of electrons that can be accommodated in each orbital = 2.
• Maximum number of electrons that can be accommodated in each subshell:
◦ s-subshell: 2 (1 orbital × 2 electrons/orbital).
◦ p-subshell: 6 (3 orbitals × 2 electrons/orbital).
◦ d-subshell: 10 (5 orbitals × 2 electrons/orbital).
◦ f-subshell: 14 (7 orbitals × 2 electrons/orbital).
Filling of Electrons in Subshells (Subshell Electron Configuration)
• Subshells are represented with the principal quantum number 'n' (e.g., 1s, 2s, 2p).
• Electrons are filled gradually in the subshells in the increasing order of energy.
• Energy Order Rule (n+l rule):
◦ The energy of subshells generally increases with the ascending order of (n+l) values.
◦ If two subshells have the same (n+l) value, the subshell with the higher 'n' value has more energy.
◦ Example: 1s (n+l=1) < 2s (n+l=2) < 2p (n+l=3) < 3s (n+l=3) < 3p (n+l=4) < 4s (n+l=4) < 3d (n+l=5).
▪ Note that 4s has lower energy than 3d, so 4s fills before 3d.
• Examples of Subshell Electron Configuration:
◦ Hydrogen (1H): 1s¹.
◦ Helium (2He): 1s².
◦ Lithium (3Li): 1s² 2s¹.
◦ Neon (10Ne): 1s² 2s² 2p⁶.
◦ Potassium (19K): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ (last electron in 4s).
◦ Scandium (21Sc): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹ (or 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹ 4s² in order of shell).
▪ In transition elements, 4s subshell fills completely first, then 3d subshell fills gradually.
• Noble Gas Configuration (Short Form)
◦ The symbol of the preceding noble gas is placed in square brackets, followed by the electron configuration of the remaining subshells.
◦ Example: Sodium (11Na): 1s² 2s² 2p⁶ 3s¹ can be written as [Ne] 3s¹.
Peculiarity of Chromium (Cr) and Copper (Cu)
• Stability of Half-filled (d⁵) and Completely Filled (d¹⁰) Configurations: These configurations are more stable than other d-configurations (d¹ to d¹⁰).
• Chromium (24Cr): Expected: 3d⁴ 4s²; Actual stable: 3d⁵ 4s¹. (One electron from 4s moves to 3d to achieve a stable half-filled d-subshell).
• Copper (29Cu): Expected: 3d⁹ 4s²; Actual stable: 3d¹⁰ 4s¹. (One electron from 4s moves to 3d to achieve a stable completely filled d-subshell).
Finding Block, Period, and Group from Subshell Electron Configuration
1. Block:
◦ The subshell to which the last electron was added determines the block.
◦ s-block: last electron in s-subshell (Groups 1, 2).
◦ p-block: last electron in p-subshell (Groups 13-18).
◦ d-block: last electron in d-subshell (Groups 3-12).
◦ f-block: last electron in f-subshell (placed separately at bottom).
2. Period Number:
◦ The highest shell number (n) in its subshell electron configuration represents the period number.
◦ Example: Carbon (6C) 1s² 2s² 2p¹ has highest shell number 2, so it's in Period 2.
3. Group Number:
◦ s-block elements: The number of electrons in the outermost s subshell.
▪ Example: Li (1s² 2s¹) has 1 electron in outermost s, so Group 1.
▪ Mg (1s² 2s² 2p⁶ 3s²) has 2 electrons in outermost s, so Group 2.
◦ p-block elements: 10 + (number of electrons in the outermost s and p subshells). (The '10' accounts for the 10 d-block groups).
▪ Example: Silicon (14Si) 1s² 2s² 2p⁶ 3s² 3p² has 2+2=4 outermost electrons, so Group 10+4 = 14.
◦ d-block elements: The sum of the number of electrons in the outermost s subshell and the number of electrons in the d subshell preceding it.
▪ Example: Scandium (21Sc) 3d¹ 4s² has 1+2=3 electrons, so Group 3.
▪ Iron (26Fe) 3d⁶ 4s² has 6+2=8 electrons, so Group 8.
Periodic Trends in Periodic Table
Ionisation Enthalpy
• Definition: The minimum amount of energy required to remove the most loosely bound electron from the outermost shell of an isolated gaseous atom.
• Down a Group: Decreases.
◦ Number of shells increases, increasing the distance of outermost electrons from the nucleus.
◦ Although nuclear charge increases, its influence is overcome by the increasing number of shells. This makes it easier to remove the outermost electron.
• Across a Period (Left to Right): Increases.
◦ No change in the number of shells.
◦ Nuclear charge gradually increases, leading to a stronger attractive force of the nucleus on the outermost electrons. This makes it harder to remove the outermost electron.
• Location: Elements with low ionisation enthalpy are generally on the left side (s-block). Elements with high ionisation enthalpy are generally on the right side (p-block, especially noble gases). Caesium and Francium have the least ionisation enthalpy.
Characteristics of s-block Elements
• Includes alkali metals (Group 1) and alkaline earth metals (Group 2).
• Valence shell electron configurations are ns¹ (Group 1) and ns² (Group 2). Hydrogen (1s¹) and Helium (1s²) are s-block elements.
• When they participate in chemical reactions, they generally donate electrons from their outermost s subshell.
• Exhibit fixed oxidation states: +1 for Group 1 and +2 for Group 2.
• Generally exist in the solid state. Caesium (Cs) has a very low melting point and can be liquid on warm days. Francium (Fr) and Radium (Ra) are radioactive.
• Radium (Ra): An alkaline earth metal, radioactive, discovered by Marie and Pierre Curie. Used in early cancer treatment, now restricted due to health risks.
Characteristics of p-block Elements
• Includes Groups 13 to 18.
• Comprise metals, non-metals, and metalloids, existing in solid, liquid, and gaseous states.
• Exhibit both positive (+) and negative (–) oxidation states.
• Gallium (Ga) has a very low melting point (29.77°C) and can be liquid on warm days.
• Along with s-block elements, they are considered main group elements.
Characteristics of d-block Elements (Transition Elements)
• Placed in Groups 3 to 12 of the periodic table.
• Electrons are gradually filled into the penultimate (n-1)d subshell.
• All d-block elements are metals.
• Similarities in Properties: Show similarities not only within their groups but also along the periods, because their outermost subshell electron configuration (ns¹⁻²) is generally the same.
• Variable Oxidation States:
◦ They show variable oxidation states (e.g., Iron (Fe) can be +2 in FeCl2 and +3 in FeCl3).
◦ This is because there is only a slight energy difference between the outermost s subshell and the penultimate d subshell. Under favorable conditions, electrons from the d subshell also participate in chemical reactions, in addition to s electrons.
• Coloured Compounds: Compounds of transition elements are generally coloured (e.g., copper sulphate, potassium permanganate, potassium dichromate). This is due to the presence of their ions or ions containing them.
◦ Zinc (Zn) compounds are colorless as it is a pseudo transition element.
• Pseudo Transition Elements: Zinc (Zn), Cadmium (Cd), and Mercury (Hg) (Group 12) do not show all the general characteristics of transition elements because their d subshell is completely filled.
• Titanium (Ti): A strong, lightweight metal with low density and high melting point. It does not corrode easily. Used in aircraft, spacecraft, missiles, medical implants (non-toxic, non-allergic), and paint manufacturing (titanium dioxide).
Characteristics of f-block Elements (Inner Transition Elements)
• Located at the bottom of the periodic table in two separate rows.
• Electrons are filled in the anti-penultimate shell.
• Also known as inner transition elements.
• Lanthanoids: f-block elements of the 6th period.
• Actinoids: f-block elements of the 7th period.
• Properties:
◦ Show variable oxidation states.
◦ Actinoids are radioactive, and many are man-made.
• Uses:
◦ Isotopes of Uranium (U), Thorium (Th), and Plutonium (Pu) are used as fuel in nuclear reactors.
◦ Neodymium (Nd) is used for strong magnets.
◦ Cerium (Ce) and Lanthanum (La) are used as catalysts in the petroleum industry.
• Rare Earth Elements: A group of 17 elements (15 lanthanoids plus scandium and yttrium). Though not truly rare, they are scattered, making extraction challenging. They have diverse applications in modern technology (computers, LCD screens, mobile phones, renewable energy, batteries). Monazite, an important ore, is found in coastal Kerala.
--------------------------------------------------------------------------------
Unit 4: Gas Laws and Mole Concept
This unit explores the general properties of gases and fundamental concepts in chemistry related to quantity of matter.
Introduction to Gases
• Gases have the least density among the three states of matter (solid, liquid, gas).
• Arrangement of Particles in Gaseous State:
◦ Distance between molecules: Very high.
◦ Force of attraction between molecules: Very low.
◦ Freedom of movement of molecules: Very high.
◦ Energy of molecules: High.
• Kinetic Molecular Theory (James Clerk Maxwell and Ludwig Boltzmann)
◦ Gases consist of minute particles (atoms/molecules).
◦ Attractive force between gaseous molecules is very low.
◦ Volume of gaseous molecules is negligible compared to total gas volume.
◦ Volume can be reduced by reducing distance between molecules.
◦ Gaseous molecules are in constant, random motion, colliding with each other and container walls. These collisions with container walls result in gaseous pressure.
◦ Collisions are elastic (kinetic energy before and after collision is the same).
◦ Average kinetic energy of gaseous molecules is directly proportional to its temperature.
General Properties of Gases
1. Volume (V)
◦ The space occupied by a substance.
◦ The volume of a gas is the volume of the container in which it is occupied.
◦ Units: Litre (L), cubic centimeter (cm³ or cc), milliliter (mL), cubic meter (m³).
▪ 1 cm³ = 1 mL.
▪ 1000 cm³ = 1000 mL = 1 L.
▪ SI unit: m³.
▪ 1 m³ = 1000 L.
2. Pressure (P)
◦ Definition: The force experienced per unit surface area due to collisions of gaseous molecules with container walls. P = Force (F) / Surface area (A).
◦ Units: Atmospheric pressure (atm), Pascal (Pa), Newton per square meter (N/m²).
▪ 1 atm = 1.01325 × 10⁵ Pa.
◦ Measured using a manometer.
3. Temperature (T)
◦ Related to the kinetic energy of molecules; heating a gas increases molecular kinetic energy and thus temperature.
◦ Units: Kelvin (K) is the SI unit; Celsius (°C) is a common unit.
◦ Conversion: To convert °C to Kelvin, add 273: t°C = (t + 273) K.
Gas Laws
1. Boyle’s Law (Pressure-Volume Relation)
◦ Proposed by Robert Boyle in 1662.
◦ Statement: At constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure.
◦ Mathematical Representation: V α 1/P (temperature, mass constant) or PV = k (where k is a constant).
◦ For two states: P₁V₁ = P₂V₂.
◦ Graph: A pressure-volume (P-V) graph shows an inverse relationship, and a PV vs P graph shows a constant value.
◦ Applications: Weather balloons expanding at higher altitudes, air bubbles growing as they rise in water.
2. Charles’s Law (Volume-Temperature Relation)
◦ Studied by Jacques Alexandre Ceasre Charles.
◦ Absolute Zero: Extrapolating the volume-temperature graph shows that the volume of a gas becomes zero at −273.15°C (or −273°C for practical purposes). This temperature is called Absolute Zero and is the basis of the Kelvin scale (0 K).
◦ Statement: At constant pressure, the volume of a definite mass of a gas is directly proportional to its temperature in Kelvin Scale.
◦ Mathematical Representation: V α T (pressure, mass constant) or V/T = k (where k is a constant).
◦ For two states: V₁/T₁ = V₂/T₂.
◦ Applications: Filling vehicle tyres with lower air pressure in summer, submerging liquid ammonia containers in cold water before opening.
3. Avogadro’s Law (Volume-Number of Particles Relation)
◦ Statement: At constant temperature and pressure, equal volumes of all gases contain an equal number of molecules. Conversely, equal numbers of molecules of different gases occupy equal volumes.
◦ Mathematical Representation: V α N (temperature, pressure constant), where N is the number of molecules.
◦ Applications: Inflating balloons, filling air in footballs.
Combined Gas Equation
• Combines Boyle's and Charles's laws:
◦ V α 1/P (T, n constant)
◦ V α T (P, n constant)
• Resulting Equation: (P₁V₁)/T₁ = (P₂V₂)/T₂ for a definite mass of gas when pressure, volume, and temperature change.
Ideal Gas Equation
• Combines Boyle's, Charles's, and Avogadro's laws:
◦ V α 1/P
◦ V α T
◦ V α n (n = number of moles)
• Resulting Equation: PV = nRT (where R is the Universal Gas Constant).
• Ideal Gases: Gases that obey this equation at all temperatures and pressures.
Mole Concept
• Definition: A mole is the quantity of a substance containing 6.022 × 10²³ particles (atoms/molecules/ions).
• Avogadro Number (NA): This specific number (6.022 × 10²³) is called Avogadro number, named after Amedeo Avogadro. Its value has been refined over time.
• Significance: The mole is the SI unit of quantity of matter. It enables accurate measurement of reactants and products in chemical reactions.
Relative Atomic Mass and Mole
• Relative Atomic Mass: Expresses the mass of one atom relative to another. It indicates how many times an atom is heavier compared to 1/12th the mass of a carbon-12 atom, which is considered a single unit (unified mass, u or Da).
• Average Atomic Mass: Atomic masses often have decimal values because they are calculated as the average mass of an element's isotopes based on their natural abundance.
◦ Example: Chlorine's average atomic mass is 35.5 u due to 35Cl (75%) and 37Cl (25%).
• Gram Atomic Mass: An element that weighs as much as its relative atomic mass in grams contains 6.022 × 10²³ atoms. This mass is known as gram atomic mass. One gram atomic mass contains 1 mole of atoms.
Molar Mass
• Molecular Mass: The total mass of atoms in a molecule.
◦ Example: CO₂ molecular mass = 1(C) + 2(O) = 1(12) + 2(16) = 44.
• Gram Molecular Mass (Molar Mass): The molecular mass of a substance expressed in grams. One molar mass of a compound contains one mole of molecules.
◦ Example: 44g of CO₂ is 1 mole of CO₂ molecules.
• Calculating Moles: Number of moles = Given mass / Molar mass.
◦ Number of molecules = Number of moles × 6.022 × 10²³.
Volume of Gases and Mole
• STP (Standard Temperature and Pressure): Defined as 273 K (0°C) and 1 atm pressure.
• Molar Volume at STP: At STP, one mole of any gas occupies a volume of 22.4 L.
• Calculating Moles from Volume (at STP): Number of moles of gases at STP = Given volume (in litres) / Molar volume at STP (22.4 L).
Mole Concept and Chemical Equations
• Balanced chemical equations represent mole ratios of reactants and products.
◦ Example: 2H₂ + O₂ → 2H₂O.
▪ This means 2 moles of H₂ react with 1 mole of O₂ to form 2 moles of H₂O.
• These mole ratios can be converted to mass ratios using molar masses:
◦ Mass of 2 mol H₂ = 2 × (2g/mol) = 4g.
◦ Mass of 1 mol O₂ = 32g.
◦ Mass of 2 mol H₂O = 2 × (18g/mol) = 36g.
◦ So, 4g H₂ + 32g O₂ → 36g H₂O.
• This allows for stoichiometric calculations to determine the amount of reactants needed or products formed.
◦ Example: If 40g H₂ reacts, 360g H₂O is formed, requiring 320g O₂.
• Chemical equations are vital for controlling processes like neutralising acid wastes in industries, by calculating the required amounts of substances like lime or baking soda.
