2Language of Chemistry 2nd version

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MATTER
▢ Definition: Fixed mass + occupies definite space + felt by senses → matter
▢ Universe:

Table 1: Universe → Energy + Matter

Universe
Primary divisions
Classification
Energy
—
Matter
Physical → solid / liquid / gas
Chemical → pure substance / mixture
▢ Matter–Energy: Interconvertible by Einstein mass–energy equation: \(E = mc^2\)
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CLASSIFICATION OF MATTER

Table 1: Physical + Chemical Classification

Basis
Classes
Subclasses
Physical state
Solid
Definite shape + definite volume
Liquid
Definite volume + no definite shape
Gas
No definite shape + no definite volume
Chemical purity
Pure substances
Element → metals / non-metals / metalloids; Compound → organic / inorganic
Impure substances / Mixture
Homogeneous / Heterogeneous
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STATES OF MATTER
📖
Solid
📝
Features
  • •
    Definite shape + volume + mass
  • •
    Hard; strong attraction between constituent particles
  • •
    Incompressible; examples: salt, wood, stone
  • •
    Particles held by strong forces: ionic bond / covalent bond / van der Waals force
  • •
    Atoms / ions / molecules → vibrate around fixed positions
  • •
    Rigid
📖
Liquid
📝
Features
  • •
    Definite mass + volume; no definite shape → takes vessel shape
  • •
    Universal character: volume conservation
  • •
    Interparticle attraction: not very high
  • •
    Compressibility: very low
  • •
    Molecules can translate
  • •
    Examples: water, alcohol, milk, oil
Ideal Liquid
Ideal liquid is incompressible. Pascal's law.
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Gas
📝
Features
  • •
    Definite mass; no definite shape + volume
  • •
    Intermolecular attraction very low; generally weak van der Waals force
  • •
    Intermolecular space large
  • •
    Highly compressible
  • •
    High rate of diffusion
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ELEMENT
▢ Concept: Introduced by Robert Boyle
▢ Definition: Simplest substance → cannot be decomposed into simpler substances by ordinary physical / chemical processes
▢ Number:
  • •
    Known elements ≈ 114
  • •
    Natural ≈ 92
  • •
    Artificial ≈ 22
▢ Artificial Elements: Synthesized by special nuclear processes.
▢ Occurrence:
  • •
    ≈ \(\frac{1}{4}\) natural elements → free state
  • •
    Other elements → combined state
▢ Examples:
❖ Free existing: \(H_2, O_2, N_2, Au, Pt\)
❖ Combined state: \(Na, Ca, Zn\)
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CLASSIFICATION OF ELEMENTS
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Metals
  • •
    Electropositive character
  • •
    Usually solid; exception: \(Hg\)
  • •
    Malleable + lustrous + ductile
  • •
    High melting point + boiling point
  • •
    Good conductors of heat + electricity; examples: \(Ag, Cu, Au, Fe, Na\)
  • •
    Low ionisation energy
  • •
    Oxides generally basic
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Non-metals
  • •
    Electronegative character
  • •
    Exist in all three states; \(C\) solid, \(Br_2\) liquid, \(Cl_2\) gas
  • •
    Neither malleable nor ductile
  • •
    Non-lustrous; exception: iodine
  • •
    Low melting point + boiling point; exception: diamond
  • •
    Bad conductors of heat + electricity; exception: graphite
  • •
    High ionisation energy
  • •
    Oxides generally acidic
📖
Metalloids
📝
Definition
Elements showing intermediate characters between metals and non-metals
📝
Examples
\(As, C, Sb, Ge, Te, Ga\)
Graphite
Graphite has free electrons due to \(sp^2\) hybridization + planar structure.
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COMPOUND
▢ Definition: Two or more elements + definite proportion by weight → compound
▢ Properties:
  • •
    Properties different from constituent elements
  • •
    Homogeneous
  • •
    Decomposed into elements only by chemical process
▢ Examples: \(H_2O, CO_2, NaCl\)
▢ Inorganic Compounds:
  • •
    Obtained from non-living sources / minerals
  • •
    Generally soluble in water
  • •
    Generally insoluble in organic solvents
  • •
    Ionic reactions → fast
  • •
    Example: \(NaCl + AgNO_3 \rightarrow AgCl + NaNO_3\)
▢ Organic Compounds:
  • •
    Obtained from living plants / animals
  • •
    Usually carbon containing with \(H_2, O_2, N_2, X, S\)
  • •
    Generally insoluble in water
  • •
    Soluble in organic solvents
  • •
    Low melting point + boiling point
  • •
    Volatile
  • •
    Molecular reactions → slow
  • •
    Examples: sugar, alcohol, vitamins, enzymes
Organic Compounds Soluble in Water
Some organic compounds dissolve in water due to intermolecular H-bonding; examples: alcohols, acids.
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MIXTURE
▢ Definition: Two or more substances in any weight proportion; components retain identity → mixture
▢ Separation: Components separated by physical methods.
▢ Types:

Table 1: Types of Mixture

Type
Key feature
Examples
Heterogeneous mixture
Two or more phases; components not uniformly distributed
sand + water; rice + gravels
Homogeneous mixture
One phase; components uniformly distributed
milk + water; alcohol + water
▢ Special Points:
  • •
    Compound → always homogeneous
  • •
    Mixture → homogeneous or heterogeneous
  • •
    Mixture separation → physical / chemical methods
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ATOM, MOLECULE, VALENCY, REAGENT
📖
Atom
📝
Definition
Smallest particle of element taking part in chemical reaction; may or may not have independent existence
📝
Term coined by
Dalton
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Molecule
📝
Definition
Smallest particle of element / compound having all properties + free existence
📝
Term introduced by
Avogadro
📝
Types
📄
Homoatomic molecules
Similar atoms only; examples: noble gases, \(H_2, F_2, O_3, P_4, S_8\)
📄
Heteroatomic molecules
Different atoms; examples: \(HCl, NH_3, CO_2\)
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Valency
📝
Old concept
Number of \(H\)-atoms or \(Cl\)-atoms with which one atom combines
📝
Modern concept
Number of electrons lost / gained / shared by one atom → stable electronic configuration of nearest inert gas
📖
Substrate
Substance on which reagent attacks; generally high molecular weight.
📖
Reagent
Attacking species in chemical reaction; generally low molecular weight.
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Precipitate
Insoluble substance formed by interaction of two solutions; phenomenon → precipitation.
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CHEMICAL EQUATION
▢ Definition: Symbolic expression of chemical reaction.
▢ Essentials:
  • •
    Represents true chemical reaction
  • •
    Molecular form
  • •
    Balanced for exact quantitative information
▢ Example: \(2H_2 + O_2 \rightarrow 2H_2O\)
Balancing
Balancing of chemical equation is based on conservation of mass.
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RADICALS / IONS
▢ Definition: Charged particles / species → ions or radicals.
▢ Radical: Atom / group of atoms with \(+\) or \(-\) charge acting as single unit during chemical reaction.
▢ Basic Radical / Cation:
  • •
    Electropositive radical
  • •
    Positively charged
  • •
    From bases during salt formation
  • •
    Moves towards cathode during electrolysis
  • •
    Examples: \(Na^+, Ca^{2+}, Al^{3+}, NH_4^+\)
▢ Acid Radical / Anion:
  • •
    Electronegative radical
  • •
    Negatively charged
  • •
    From acids during salt formation
  • •
    Examples: \(Cl^-, S^{2-}, N^{3-}, C^{4-}, SO_4^{2-}, NO_3^-\)
▢ Simple Radical: Only one atom; examples: \(Na^+, S^{2-}, Mg^{2+}\)
▢ Compound Radical: Two or more atoms; examples: \(NH_4^+, CO_3^{2-}, SO_4^{2-}, HS^-\)

Table 1: Special Radicals

No.
Radical
Formula
1
Ammonium
\(NH_4^+\)
2
Chromium
\(Cr^{3+}\)
3
Ferrous / Ferric
\(Fe^{2+}\) / \(Fe^{3+}\)
4
Aurous / Auric
\(Au^+\) / \(Au^{3+}\)
5
Cuprous / Cupric
\(Cu^+\) / \(Cu^{2+}\)
6
Bisulphite
\(HSO_3^-\)
7
Chlorate
\(ClO_3^-\)
8
Cyanate
\(CNO^-\)
9
Perchlorate
\(ClO_4^-\)
10
Bisulphide
\(HS^-\)
11
Hypochlorite
\(ClO^-\)
12
Hypoiodite
\(IO^-\)
13
Iodate
\(IO_3^-\)
14
Meta-aluminate
\(AlO_2^-\)
15
Metaphosphate
\(PO_3^-\)
16
Permanganate
\(MnO_4^-\)
17
Manganate
\(MnO_4^{2-}\)
18
Chromate
\(CrO_4^{2-}\)
19
Dichromate
\(Cr_2O_7^{2-}\)
20
Molybdate
\(MoO_4^{2-}\)
21
Oxalate
\(C_2O_4^{2-}\)
22
Plumbate
\(PbO_2^{2-}\)
23
Thiosulphate
\(S_2O_3^{2-}\)
24
Tetrathionate
\(S_4O_6^{2-}\)
25
Zincate
\(ZnO_2^{2-}\)
26
Aluminate
\(Al_2O_3^{2-}\)
27
Ferricyanide
\([Fe(CN)*6]^{3-}\)
28
Borate
\(BO_3^{3-}\)
29
Arsenate
\(AsO_4^{3-}\)
30
Carbide
\(C^{4-}\)
31
Cobalt nitrite
\([Co(NO_2)*6]^{3-}\)
32
Ferrocyanide
\([Fe(CN)*6]^{4-}\)
33
Pyrophosphate
\(P_2O_7^{4-}\)
34
Tetra / pyroborate
\(B_2O_7^{2-}\)
35
Sulphocyanide / Thiocyanide
\(CNS^-\)
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CATALYST
▢ Definition: Substance altering reaction rate without permanent change in mass + composition.
▢ Catalysis: Phenomenon of catalyst action.
▢ Types:

Table 1: Positive vs Negative Catalyst

Type
Effect
Activation energy
Example
Positive catalyst
Increases rate
Decreases \(E_a\)
\(2KClO_3 \xrightarrow{MnO_2,\ 250^\circ C} 2KCl + 3O_2\) fast
Negative catalyst
Decreases rate
Increases \(E_a\)
\(2H_2O_2 \xrightarrow{dil.\ H_2SO_4/glycerol} 2H_2O + O_2\) slow
▢ Examples:
  • •
    \(2KClO_3 \xrightarrow{400^\circ C} 2KCl + 3O_2\) slow
  • •
    \(2KClO_3 \xrightarrow{MnO_2,\ 250^\circ C} 2KCl + 3O_2\) fast
  • •
    \(2H_2O_2 \xrightarrow{Pt} 2H_2O + O_2\) very fast
▢ Auto-catalyst:
❖ Definition: One product itself acts as catalyst.
❖ Examples:
  • •
    Acidified oxalic acid + \(KMnO_4\): \(Mn^{2+}\) acts as auto-catalyst
  • •
    \(CH_3COOC_2H_5 + H_2O \rightarrow CH_3COOH + C_2H_5OH\): acetic acid acts as auto-catalyst
▢ Characteristics:
  • •
    Only alters reaction rate
  • •
    Mass + composition unchanged at end
  • •
    Small amount sufficient
  • •
    Cannot initiate reaction
  • •
    Specific action
Extra Points
  • •
    Acid–base reaction → double displacement reaction
  • •
    Symbol → abbreviation of element name
  • •
    Sand + powdered charcoal mixture separated by washing
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SPECIAL TERMS

Table 1: Iso-terms

Term
Meaning
Examples / Notes
Isotopes
Same atomic number \(Z\); different mass number \(A\)
\(^1_1H, ^2_1H, ^3_1H\); \(^{40}*{20}Ca, ^{41}*{20}Ca, ^{42}*{20}Ca\)
Isobars
Different atomic number; same mass number / same nucleons
\(^{40}*{18}Ar, ^{40}*{19}K, ^{40}*{20}Ca\)
Isotones
Different \(Z\) + different \(A\); same neutron number
\(^{30}*{14}Si, ^{31}_{15}P\)
Isobels
Same geometry + same hybridization
\(SO_2, SnCl_2\)
Isosteres
Isoelectronic + isoatomic
\(N_2O, CO_2\)
Isomorphs
Same number of atoms bonded similarly
\(FeSO_4\cdot7H_2O, MgSO_4\cdot7H_2O, ZnSO_4\cdot7H_2O\); \(K_2SO_4, K_2CrO_4, K_2SeO_4\)
Alums
Isomorphous compounds
\(K_2SO_4\cdot Al_2(SO_4)_3\cdot24H_2O\) potash alum; \(K_2SO_4\cdot Cr_2(SO_4)_3\cdot24H_2O\) chrome alum
Isoatomic
Same number of atoms
\(N_2O, SO_2\)
Isoelectronic
Same number of electrons
\(CO, CN^-, NO^+\)
Isodiaphers
Same neutron–proton difference \((n-p)\)
\(\alpha\)-decay produces isodiapher with parent chain
Isomers
Same \(Z\) + same \(A\); different radioactive properties
\(\gamma\)-decay produces nuclear isomer
Decay Points
  • •
    \(\alpha\)-decay followed by \(\beta\)-decay → isotope of parent atom
  • •
    Mass number \(14\) + neutron number \(8\) → isotope of carbon
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ATOMIC WEIGHT / ATOMIC MASS
▢ History: Reference standard: hydrogen → oxygen-16 → carbon-12.
▢ Definition: Number showing how many times average atom of element is heavier than \(\frac{1}{12}\) mass of \(^{12}C\) atom.
▢ Formula: \(At.\ wt. = \frac{Average\ wt.\ of\ atom}{\frac{1}{12}\times mass\ of\ ^{12}C\ atom}\)
▢ Nature:
  • •
    Average relative weight
  • •
    Simple ratio; no true unit
  • •
    Expressed in amu
▢ Unit:
  • •
    \(1\ amu = \frac{1}{12}\times mass\ of\ ^{12}C\ atom\)
  • •
    Mass of one \(^{12}C\) atom \(= 1.9924\times10^{-23}\ g\)
  • •
    \(1\ amu = \frac{1.9924\times10^{-23}}{12}\ g = 1.66\times10^{-24}\ g\)
▢ Examples:
  • •
    Lightest known atom: hydrogen
  • •
    Heaviest naturally occurring atom: uranium-238
  • •
    \(At.\ mass\ of\ Na = 23\ amu\) → one Na atom = 23 times heavier than \(1\ amu\)
▢ Fractional Atomic Weight:
❖ Cause: Stable isotopes with different relative abundance
❖ Chlorine example: \(\frac{75\times35 + 25\times37}{100}=35.5\ amu\)
▢ Average Weight of One Atom: \(Average\ wt.\ of\ 1\ atom = At.\ wt.\times1.66\times10^{-24}\ g\)
▢ Gram Atomic Weight:
❖ Definition: Atomic weight expressed in gram → gram atomic weight \((GAW)\)
❖ Examples: \(1\ GAW\ Na = 23\ g\); \(1\ GAW\ Ca = 40\ g\); \(1\ GAW\ S = 32\ g\)
❖ Avogadro number: \(1\ GAW\ of\ every\ element = 6.023\times10^{23}\ atoms\)
▢ Gram Atom / Mole Atom:
❖ Definition: Atomic weight of element in grams = one gram atom / mole atom
❖ Number: \(No.\ of\ gm\ atoms=\frac{wt.\ of\ element\ in\ g}{GAW\ of\ element}\)
Composition Terms
  • •
    Gravimetric composition → by weight; water \(H:O = 1:8\)
  • •
    Volumetric composition → by volume; water \(H:O = 2:1\)
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METHODS OF DETERMINING ATOMIC WEIGHT

Table 1: Atomic Weight Methods

Method
Use / Formula
Important Points
Dulong–Petit method
\(Approx.\ At.\ wt.\times Specific\ heat = 6.4\); \(At.\ wt.\ (appx.)=\frac{6.4}{Sp.\ heat}\)
Only solid elements; exceptions: \(Be, B, C, Si\); \(Exact\ At.\ wt.=Eq.\ wt.\times Valency\); \(Valency=\frac{App.\ At.\ wt.}{Eq.\ wt.}\)
Vapour density method
\(Valency=\frac{Mol.\ wt.\ of\ chloride}{Eq.\ wt.\ of\ chloride}=\frac{2\times V.D.\ of\ chloride}{Eq.\ wt.\ of\ metal+35.5}\)
Suitable for elements with volatile chlorides; \(At.\ wt.=Eq.\ wt.\ of\ metal\times Valency\)
Specific heat method
\(At.\ wt.\ of\ gaseous\ element=\frac{Mol.\ wt.}{Atomicity}\)
Atomicity: noble gases \(1\), diatomic gases \(2\), ozone \(3\), \(P/As=4\), \(S/Se/Te=8\)
Isomorphism method
Identical crystal structure → similar constitution + chemical formula
Examples: \(K_2SO_4, K_2CrO_4, K_2SeO_4\); \(ZnSO_4\cdot7H_2O, MgSO_4\cdot7H_2O, FeSO_4\cdot7H_2O\); alums
Cannizzaro's method
\(At.\ mass=\frac{Minimum\ molecular\ mass\ of\ compound\times \%\ element}{100}\)
Used for non-metals: \(C, N, S, O\); mass spectrometer also determines molecular mass
📖
Condition for isomorphism
  • •
    Similar crystalline structures
  • •
    Same number of similar atoms
  • •
    Same total ionic charge
  • •
    Similar bonding nature
📖
Isomorphism relation
\(\frac{Wt.\ of\ A\ combines\ with\ other\ elements}{Wt.\ of\ B\ combines\ with\ other\ elements}=\frac{At.\ wt.\ of\ A}{At.\ wt.\ of\ B}\)
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MOLECULAR WEIGHT / MOLECULAR MASS
▢ Definition: Number showing how many times one molecule is heavier than \(\frac{1}{12}\) mass of \(^{12}C\) atom.
▢ Formula: \(Mol.\ wt.=\frac{wt.\ of\ one\ molecule\ of\ substance}{\frac{1}{12}\times wt.\ of\ ^{12}C}=\frac{wt.\ of\ one\ molecule}{amu}\)
▢ Nature:
  • •
    Ratio → no true unit
  • •
    Expressed in amu
  • •
    Sum of atomic masses of all atoms in molecule
▢ Example: \(H_2SO_4: Mol.\ wt.=2+32+64=98\ amu\)
▢ Gram Molecular Weight:
❖ Definition: Molecular weight expressed in gram = gram molecular weight \((GMW)\) / gram molecule / mole molecule
❖ Moles: \(No.\ of\ gm\ molecules=\frac{Wt.\ of\ substance}{GMW}\)
❖ Mass: \(Wt.\ of\ substance\ in\ g=No.\ of\ molecules\times GMW\)
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METHODS OF DETERMINING MOLECULAR WEIGHT

Table 1: Molecular Weight Methods

Method
Formula
Notes
Diffusion method / Graham's law
\(\frac{r_1}{r_2}=\sqrt{\frac{M_2}{M_1}}\)
Rate of diffusion inversely proportional to square root of molecular weight
Vapour density method
\(Mol.\ wt.=2\times Vapour\ density\)
For gases only; vapour density = relative density
Victor Meyer's method
\(Mol.\ wt.=\frac{wt.\ of\ volatile\ substance}{V_2\ ml\ of\ substance\ at\ STP}\times22400\)
For volatile liquids + solids; based on Dalton partial pressure law + Avogadro hypothesis
Berzelius hypothesis
Equal volumes of all gases contain equal number of atoms at same \(T,P\)
Contrary to Dalton's atomic theory
Avogadro's hypothesis
Equal volumes of all gases contain equal number of molecules at same \(T,P\)
Explains law of combining volumes
📖
Vapour density notes
  • •
    Molecular association → vapour density increases; examples: \(CH_3COOH, HF\)
  • •
    Dissociation → vapour density decreases; examples: \(NH_4Cl, PCl_5\)
📖
Avogadro hypothesis uses
  • •
    Deriving molecular formula of gas
  • •
    Deriving relation: \(Mol.\ wt.=2\times Vapour\ density\)
  • •
    Determining atomicity of gases
  • •
    Determining number of molecules in \(1\) gram molecular weight
  • •
    Deriving gram molecular volume
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EQUIVALENT WEIGHT
▢ Definition: Parts by weight of substance combining with / displacing \(1.008\) parts \(H\), \(8.0\) parts \(O\), or \(35.5\) parts \(Cl\).
▢ Gram Equivalent Weight: Quantity in grams numerically equal to equivalent weight.
▢ Number of Gram Equivalents: \(No.\ of\ GEW=\frac{Wt.\ of\ substance\ in\ grams}{GEW\ of\ substance}\)

Table 1: Equivalent Weight Relations

Substance
Formula
Meaning
Element
\(EW=\frac{Atomic\ weight}{Valency}\)
Valency of element
Acid
\(EW=\frac{Mol.\ wt.}{Basicity}\)
Basicity = no. of replaceable \(H^+\)
Base
\(EW=\frac{Mol.\ wt.}{Acidity}\)
Acidity = no. of replaceable \(OH^-\)
Salt
\(EW=\frac{Mol.\ wt.}{Total\ +ve\ or\ -ve\ charge}\)
Charge of cation or anion
Reducing agent
\(EW=\frac{Formula\ wt.}{No.\ of\ electrons\ lost\ per\ molecule\ or\ total\ change\ in\ O.N.}\)
Oxidation
Oxidising agent
\(EW=\frac{Formula\ wt.}{No.\ of\ electrons\ gained\ per\ molecule\ or\ total\ change\ in\ O.N.}\)
Reduction
Radical
\(EW=\frac{Formula\ weight\ of\ radical}{No.\ of\ units\ of\ charge}\)
Example: \(EW\ of\ SO_4^{2-}=\frac{96}{2}=48\)
📚
EQUIVALENT WEIGHT OF OXIDISING AGENTS

Table 1: Equivalent Weight of \(KMnO_4\) in Different Medium

Medium
Change / reaction
Equivalent weight
Acidic
\(MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O\)
\(\frac{M}{5}=\frac{158}{5}=31.60\)
Basic
\(MnO_4^- + e^- \rightarrow MnO_4^{2-}\)
\(\frac{M}{1}=158\)
Neutral
\(MnO_4^- + 4H^+ + 3e^- \rightarrow MnO_2 + 2H_2O\)
\(\frac{M}{3}=52.67\)

Table 2: Equivalent Weight of \(HNO_3\) in Different Medium

Medium
Product
Oxidation state of N
Equivalent weight
Very dilute
\(NH_3\)
\(-3\)
\(\frac{M}{8}\)
Dilute
\(N_2O\)
\(+1\)
\(\frac{M}{4}\)
Moderate \((1:1)\)
\(NO\)
\(+2\)
\(\frac{M}{3}\)
Concentrated
\(NO_2\)
\(+4\)
\(\frac{M}{1}\)
📖
Example reducing agent
📝
Oxalic acid
\(C_2O_4^{2-}\rightarrow 2CO_2+2e^-\); \(Eq.\ wt.\ of\ H_2C_2O_4=\frac{90}{2}=45\)
📚
METHODS OF DETERMINING EQUIVALENT WEIGHT

Table 1: Equivalent Weight Methods

No.
Method
Formula / Principle
1
Hydrogen displacement method
\(Eq.\ wt.\ of\ metal=\frac{Wt.\ of\ metal}{Wt.\ of\ H_2\ displaced}\times1.008\); or \(\frac{W}{V}\times11200\)
2
Oxide formation method
\(Eq.\ wt.=\frac{Wt.\ of\ metal}{Wt.\ of\ oxygen}\times8\); or \(\frac{Wt.\ of\ metal}{Vol.\ of\ O_2\ at\ STP\ in\ ml}\times5600\)
3
Chloride formation method
\(Eq.\ wt.=\frac{Wt.\ of\ metal}{Wt.\ of\ chlorine}\times35.5\); or \(\frac{Wt.\ of\ metal}{Vol.\ of\ Cl_2\ at\ STP\ in\ ml}\times11200\)
4
Neutralisation method
\(Eq.\ wt.\ of\ acid/base=\frac{W}{VN}\); \(W\)=mass in g; \(V\)=volume in L; \(N\)=normality
5
Metal displacement method
\(\frac{Wt.\ of\ metal\ added}{Wt.\ of\ metal\ displaced}=\frac{EW\ of\ metal\ added}{EW\ of\ metal\ displaced}\); more electropositive displaces less electropositive
6
Electrolytic method
\(Gram\ eq.\ wt.=Electrochemical\ equivalent\times96500\); same electricity → deposited weights ratio = equivalent weights ratio
7
Double decomposition / metathesis
\(AB+CD\rightarrow AD+CB\); \(\frac{Wt.\ AB}{Wt.\ AD}=\frac{Eq.wt.A+Eq.wt.B}{Eq.wt.A+Eq.wt.D}\); \(\frac{W_1}{W_2}=\frac{E_1}{E_2}\)
8
Conversion method
\(\frac{W_1}{W_2}=\frac{E+Eq.wt.\ of\ anion\ in\ compound\ I}{E+Eq.wt.\ of\ anion\ in\ compound\ II}\); same metal converted from compound I to compound II
9
Volatile chloride method
\(Valency=\frac{2\times VD\ of\ chloride}{E+35.5}\); \(E=\frac{2VD}{Valency}-35.5\)
10
Silver salt method
\(Eq.wt.\ of\ acid\ (RCOOH)=108\times\frac{Wt.\ of\ silver\ salt\ (RCOOAg)}{Wt.\ of\ Ag}-107\); \(Mol.\ wt.\ of\ acid=Eq.\ wt.\ of\ acid\times Basicity\)
📚
LAW OF EQUIVALENT
▢ Statement: Substances react with each other by weight in ratio of their equivalent weights.
▢ Meaning: One gram equivalent of a substance reacts completely with one gram equivalent of another substance.
▢ Examples:
  • •
    \(EW\ of\ CaCO_3=50\)
  • •
    \(EW\ of\ HCl=36.5\)
  • •
    \(50\ g\ CaCO_3\) reacts completely with \(1\ gm\ eq.\ HCl\ (36.5\ g)\)
📚
MOLE CONCEPT
▢ Avogadro Number:
❖ Symbol: \(N_A\)
❖ Value: \(6.023\times10^{23}\)
❖ Definition: Number of atoms in gram atomic mass of element OR number of molecules in gram molecular mass of substance
▢ Mole:
  • •
    Collection of \(6.023\times10^{23}\) particles
  • •
    \(1\ mole\ of\ compound = 1\ GMW\)
  • •
    \(1\ mole\ H_2O = 18\ g\)
  • •
    \(1\ mole = 6.023\times10^{23}\ molecules\)
  • •
    \(1\ mole\ gas = 22.4\ L\) at NTP
  • •
    \(1\ mole = 1\ gram\ molecular\ mass\)
▢ Main Formulas:
  • •
    \(No.\ of\ moles=\frac{given\ no.\ of\ molecules}{6.023\times10^{23}}\)
  • •
    \(No.\ of\ moles=\frac{mass\ in\ g}{Molar\ mass}\)
  • •
    \(No.\ of\ moles=\frac{given\ volume\ in\ litre\ at\ NTP}{22.4}\)
  • •
    \(No.\ of\ moles=\frac{given\ volume\ in\ ml\ at\ NTP}{22400}\)
  • •
    \(No.\ of\ moles=M\times V\), where \(M\)=molarity, \(V\)=volume in litre
  • •
    \(No.\ of\ atoms=Atomicity\times No.\ of\ molecules\)
▢ Gram Molar Volume:
❖ Definition: Volume occupied by one mole of any gas at STP.
❖ Value: \(22.4\ L\)
❖ Examples:
  • •
    \(1\ mole\ H_2 = 2\ g = 22.4\ L\)
  • •
    \(16\ g\ O_2=\frac{1}{2}\ mole=11.2\ L\)
  • •
    \(7\ g\ N_2=\frac{1}{4}\ mole=5.6\ L\)
  • •
    \(0.5\ g\ H_2=\frac{1}{4}\ mole=5.6\ L\)
  • •
    \(16\ g\ CH_4=1\ mole=22.4\ L\)
Water Exception
\(1\ mole\ H_2O\neq22400\ ml\) because ordinary \(H_2O\) is liquid; \(1\ mole\ H_2O=18\ g\) at \(273\ K\).
Points to Remember
  • •
    Mass of \(22.4\ L\ O_2\) at NTP ≈ \(32\ g\)
  • •
    Weight of \(11.2\ L\) gas at STP = vapour density in grams
  • •
    \(Density\ of\ gas\ at\ NTP=\frac{Mol.\ wt.\ in\ g}{22400\ ml}\)
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STOICHIOMETRIC CALCULATIONS
▢ Limiting Reactant / Reagent: Reactant used up first in reaction → determines amount of product formed.
▢ Example: \(MgO + H_2SO_4 \rightarrow MgSO_4 + H_2O\); \(40g\ MgO\) requires \(98g\ H_2SO_4\); if only \(90g\ H_2SO_4\), then \(H_2SO_4\) is limiting reagent.
▢ Percentage Composition:
❖ Definition: Mass percentage of each element in compound.
❖ Example: \(H_2SO_4\): \(\%S=\frac{32}{98}\times100=32.6\%\), \(\%H=\frac{2}{98}\times100=2.04\%\), \(\%O=\frac{64}{98}\times100=65.3\%\)
▢ Useful Numerical Patterns:
  • •
    \(0.224\ L\ H_2\) at STP → \(0.01\ mole\)
  • •
    \(13.5\ g\ SO_2Cl_2\); molar mass \(=135\) → moles \(=0.1\)
  • •
    Air with \(21\%\ O_2\) by volume: \(1\ L\ air\) at STP → \(210\ ml\ O_2\) → \(0.0093\ mole\)
  • •
    \(1.6\ g\ CH_4\) → \(6.023\times10^{22}\) molecules; electrons \(=10\times6.023\times10^{22}=6.023\times10^{23}\)
  • •
    Volatile substance: \(0.24\ g\) gives \(45\ ml\) vapour at NTP → mol wt \(=119.46\), \(VD=59.73\)
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EMPIRICAL FORMULA
▢ Definition: Simplest relative number of atoms of each element in compound.
▢ Determination Steps:
  1. Divide % composition / mass of each element by atomic weight → relative moles
  2. Divide all relative moles by smallest value
  3. Multiply by suitable number to make whole numbers
  4. Write empirical formula from simplest whole-number ratio
▢ Examples:
  • •
    Oxide of sulphur containing \(50\%\ S\): ratio \(S:O=\frac{50}{32}:\frac{50}{16}=1:2\) → empirical formula \(SO_2\)
  • •
    Compound of \(P\) and \(S\): \(P:S=\frac{1.24}{31}:\frac{0.96}{32}=4:3\) → empirical formula \(P_4S_3\)
  • •
    Compound with \(Na=0.0887\ mol, O=0.132\ mol, C=0.044\ mol\): ratio \(2:3:1\) → empirical formula \(Na_2CO_3\)
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MOLECULAR FORMULA
📖
_*c
📝
Definition
Actual number of atoms of each element in one molecule.
📝
Relation
\(Molecular\ formula=Empirical\ formula\times n\)
📝
Multiplier
\(n=\frac{Mol.\ wt.}{Empirical\ formula\ wt.}\)
📝
Combustion Example
1 volume gaseous compound + \(2\) volumes \(O_2\) → \(2\) volumes \(CO_2\) + \(1\) volume \(N_2\); molecular formula → \(C_2N_2\).
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DALTON'S ATOMIC THEORY

Table 1: Theory + Modification

Postulate
Modification
Matter consists of extremely small indivisible particles called atoms.
Atom has complex structure; contains electron, proton, neutron.
Atoms of an element are identical in all respects.
Isotopes have different mass number.
Atoms of different elements are different from each other.
Isobars may have same mass number.
Atoms can neither be created nor destroyed nor transformed into atoms of other elements.
Artificial transmutation can transform atoms.
Atoms combine in simple whole-number ratio to form compound atoms.
Combination may occur in complex number / ratio; example: \(C*{12}H*{22}O*{11}\).
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LAWS OF CHEMICAL COMBINATION / STOICHIOMETRY

Table 1: Main Laws

Law
Proposed / verified by
Statement
Examples / Notes
Law of conservation of mass
Lavoisier; verified by Landolt
In chemical reaction, total mass of products = total mass of reactants; mass neither created nor destroyed.
Chemical equations based on this law; not good for nuclear reactions due to \(E=mc^2\).
Law of definite proportions
Proust; verified by Stas + Richard
Pure compound always contains same elements in fixed weight ratio, irrespective of preparation method.
\(H_2O\): \(H:O=1:8\) by weight; not valid for non-stoichiometric / Berthollide compounds: \(ZnO*{0.998}, CuS, FeS, TiO\).
Law of multiple proportions
Dalton; verified by Berzelius
Different weights of one element combining with fixed weight of another bear simple numerical ratio.
\(CO\) and \(CO_2\): oxygen combining with \(12g\ C\) = \(16g\) and \(32g\) → ratio \(1:2\); examples: \(H_2O/H_2O_2\), \(SO_2/SO_3\), \(SnCl_2/SnCl_4\), oxides of nitrogen.
Law of reciprocal proportions
Richter; verified by Stas
When two elements combine separately with same weight of third element, their ratio is same or simple multiple of ratio in which they combine with each other.
Examples: \(NaCl, NaH, HCl\); \(CH_4, CO_2, H_2O\); \(NH_3, H_2O, N_2O_3\); gives law of equivalent proportions.
Gay-Lussac's law of gaseous volume
Gay-Lussac
Reacting gases combine in simple volume ratio at constant temperature + pressure.
\(H_2 + Cl_2 \rightarrow 2HCl\); volume ratio \(1:1:2\); applies only to gases; explained by Avogadro hypothesis.
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READ & DIGEST
▢ Facts:
  • •
    Loschmidt number = number of particles in \(1\ cm^3\) gas at STP = \(2.688\times10^{19}\)
  • •
    Exchange of ions by two compounds in solution → metathesis / double decomposition
  • •
    \(X_A=\frac{n_A}{n_A+n_B}\) = mole fraction of component \(A\)
  • •
    Four \(1\ L\) flasks with \(H_2, He, O_2, O_3\) at same \(T,P\): total atoms ratio \(2:1:2:3\); molecules ratio \(1:1:1:1\)
  • •
    \(Eq.\ wt.\ of\ H_2O_2=\frac{34}{2}=17\)
  • •
    \(Eq.\ wt.\ of\ K_2Cr_2O_7\) in acidic medium \(=\frac{M}{6}\), \(M=294\)
  • •
    \(Cr_2O_7^{2-}+H^+ +6e^- \rightarrow 2Cr^{3+}\)