1Language of Chemistry

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BASIC CHEMISTRY
▢ Matter and Energy:

Table 1: Matter & Energy

Term
Meaning
Matter
Substance having mass + occupying space + perceived by senses
Energy
Capacity to do work
Matter–energy relation
Interconvertible
Einstein equation
\(E = mc^2\)
▢ States of Matter:

Table 1: Solid vs Liquid vs Gas

Property
Solid
Liquid
Gas
Mass
Fixed
Fixed
Fixed
Volume
Fixed
Fixed
Variable
Shape
Fixed
Variable
Variable
Intermolecular space
Very less
Less
Large
Compressibility
Almost incompressible
Slightly compressible
Highly compressible
Intermolecular force
Very high
Moderate
Weak
Particle movement
Vibration at fixed position
Translation possible
Free random motion
Rigidity
Rigid
Fluid
Fluid
Diffusion
Nil / very slow
Low
High
Examples
Wood
Water, alcohol, milk, oil
\(H_2\), He, \(O_2\)
❖ Forces in Solids:
  • •
    Ionic bond
  • •
    Covalent bond
  • •
    Van der Waals force
❖ Forces in Liquids: Generally Van der Waals force
▢ Element:
❖ Definition: Simple substance which cannot be decomposed into or composed from simpler substances by ordinary physical / chemical process
❖ Concept Introduced By: Robert Boyle
❖ Number: ≈ 114 elements in given note → 92 natural + 22 artificial
❖ Artificial Elements: Synthesized by special nuclear processes
❖ Natural Occurrence:
  • •
    About \(\frac{1}{4}\) natural elements occur free
  • •
    Remaining occur in combined state
❖ Examples:

Table 1: Element Occurrence

State
Examples
Free state
\(H_2\), \(O_2\), \(N_2\), Au, Pt
Combined state
Na, Ca, Zn
❖ Metals, Non-metals & Metalloids:

Table 1: Metals vs Non-metals vs Metalloids

Property
Metals
Non-metals
Metalloids
Nature
Electropositive
Electronegative
Intermediate characters
State
Mostly solid except Hg
All 3 states
Mostly solid
Malleability / ductility
Malleable + ductile
Non-malleable + non-ductile
Variable
Lustre
Lustrous
Non-lustrous except iodine
Sometimes lustrous
Melting / boiling point
High
Low
Intermediate
Conduction
Good conductor of heat + electricity
Poor conductor
Semiconducting
Ionization energy
Low
High
Intermediate
Oxides
Generally basic
Generally acidic
Often amphoteric
Examples
Ag, Cu, Fe, Na
C, \(Br_2\), \(Cl_2\)
As, Sb, Ge, Te, Ga
▢ Compound:
❖ Definition: Substance formed by combination of 2 or more elements in definite proportion by weight
❖ Properties:
  • •
    Properties differ from constituent elements
  • •
    Homogeneous substance
  • •
    Can be analysed into elements by chemical process only
❖ Examples: \(H_2O\), \(CO_2\), NaCl
❖ Inorganic Compound:

Table 1: Inorganic Compounds

Feature
Point
Source
Non-living sources / minerals
Solubility
Soluble in water; insoluble in organic solvents
Reaction type
Ionic reactions
Reaction speed
Fast
Example
\(NaCl + AgNO_3 \rightarrow AgCl + NaNO_3\)
❖ Organic Compound:

Table 1: Organic Compounds

Feature
Point
Source
Living plants + animals
Elements
C, H, O, N, halogens, S etc.
Solubility
Usually insoluble in water; soluble in organic solvents
Melting / boiling point
Low
Volatility
Volatile
Reaction type
Molecular reactions
Reaction speed
Slow
Examples
Sugar, alcohol, vitamins, enzymes
◉ Exception: Alcohols + acids may dissolve in water due to intermolecular H-bonding
▢ Mixture:
❖ Definition: Collection of 2 or more substances in any proportion by weight; components retain identity
❖ Separation: Components separated by physical methods
❖ Types:

Table 1: Types of Mixtures

Type
Feature
Examples
Heterogeneous mixture
2 or more phases; components not uniformly distributed
Sand + water; rice + gravels
Homogeneous mixture
Single phase; components uniformly distributed
Milk + water; alcohol + water
❖ Special Points:
  • •
    Compound is always homogeneous
  • •
    Mixture may be homogeneous or heterogeneous
  • •
    Mixture components separated by physical / chemical methods
  • •
    Atmolysis separates gases / vapours by different diffusion rates through porous substance
❖ Atmolysis Examples:
  • •
    \(H_2 + O_2\)
  • •
    \(CH_4 + O_2\)
  • •
    \(U^{235} + U^{237}\)
▢ Atom:
❖ Definition: Smallest particle of element taking part in chemical reaction; may or may not exist independently
❖ Term Coined By: Dalton
▢ Molecule:
❖ Definition: Smallest particle of element / compound having all properties and free existence
❖ Term Coined By: Avogadro

Table 1: Homoatomic vs Heteroatomic Molecules

Type
Atoms
Examples
Homoatomic molecule
Similar atoms only
Noble gases, \(H_2\), \(F_2\), \(O_3\), \(P_4\), \(S_8\)
Heteroatomic molecule
Different atoms
HCl, \(NH_3\), \(CO_2\)
▢ Valency:
❖ Old Concept: Combining capacity of an element
❖ Modern Concept: Number of electrons lost, gained or shared by one atom to attain stable configuration of nearest inert gas
▢ Substrate & Reagent:

Table 1: Substrate vs Reagent

Term
Meaning
General molecular weight
Substrate
Substance on which reagent attacks
High
Reagent
Attacking species in chemical reaction
Low
▢ Precipitate:
❖ Definition: Insoluble substance formed by interaction of 2 solutions
❖ Phenomenon: Precipitation
▢ Chemical Equation:
❖ Definition: Symbolic expression of chemical reaction
❖ Essentials:
  • •
    Represents true chemical reaction
  • •
    Molecular
  • •
    Balanced for exact quantitative information
❖ Example: \(2H_2 + O_2 \rightarrow 2H_2O\)
❖ Balancing Principle: Conservation of mass
▢ Radicals / Ions:
❖ Definition: Charged particles / species
❖ Radical: Atom or group of atoms having +ve or -ve charge and acting as single unit in chemical reaction
❖ Types:

Table 1: Types of Radicals

Type
Charge
Other name
Origin in salt formation
Electrolysis movement
Examples
Electropositive radical
+ve
Cation / basic radical
From base
Toward cathode
\(Na^+\), \(Ca^{2+}\), \(Al^{3+}\), \(NH_4^+\)
Electronegative radical
-ve
Anion / acid radical
From acid
Toward anode
\(Cl^-\), \(S^{2-}\), \(N^{3-}\), \(SO_4^{2-}\), \(NO_3^-\)
❖ Simple vs Compound:

Table 1: Simple vs Compound Radicals

Type
Feature
Examples
Simple radical
Single atom
\(Na^+\), \(S^{2-}\), \(Mg^{2+}\)
Compound radical
2 or more atoms
\(NH_4^+\), \(CO_3^{2-}\), \(SO_4^{2-}\), \(HS^-\)
❖ Important Basic Radicals:

Table 1: Positive / Basic Radicals

Valency
Radicals
+1
\(H^+\), \(Li^+\), \(Na^+\), \(K^+\), \(Rb^+\), \(Cs^+\), \(Ag^+\), \(Cu^+\), \(Au^+\), \(NH_4^+\), \(Hg^+\)
+2
\(Be^{2+}\), \(Mg^{2+}\), \(Ca^{2+}\), \(Sr^{2+}\), \(Ba^{2+}\), \(Zn^{2+}\), \(Cu^{2+}\), \(Fe^{2+}\), \(Mn^{2+}\), \(Cr^{2+}\), \(Co^{2+}\), \(Ni^{2+}\), \(Sn^{2+}\), \(Pb^{2+}\), \(Pt^{2+}\), \(Hg^{2+}\), \(Cd^{2+}\)
+3
\(B^{3+}\), \(Al^{3+}\), \(Fe^{3+}\), \(Cr^{3+}\), \(Mn^{3+}\), \(As^{3+}\), \(Sb^{3+}\), \(Bi^{3+}\), \(Au^{3+}\)
+4
\(C^{4+}\), \(Si^{4+}\), \(Sn^{4+}\), \(Pb^{4+}\), \(Pt^{4+}\), \(Ni^{4+}\)
+5
\(As^{5+}\), \(Sb^{5+}\)
❖ Important Acid Radicals:

Table 1: Negative / Acid Radicals

Valency
Radicals
-1
\(H^-\), \(F^-\), \(Cl^-\), \(Br^-\), \(I^-\), \(OH^-\), \(NO_2^-\), \(NO_3^-\), \(BO_2^-\), \(CN^-\), \(NC^-\), \(CNO^-\), \(CNS^-\), \(SCN^-\), \(HCO_3^-\), \(CH_3COO^-\), \(ClO^-\), \(ClO_2^-\), \(ClO_3^-\), \(ClO_4^-\), \(MnO_4^-\), \(IO^-\), \(IO_2^-\), \(IO_3^-\), \(IO_4^-\)
-2
\(O^{2-}\), \(O_2^{2-}\), \(S^{2-}\), \(CO_3^{2-}\), \(C_2O_4^{2-}\), \(SO_3^{2-}\), \(SO_4^{2-}\), \(S_2O_3^{2-}\), \(S_4O_6^{2-}\), \(CrO_4^{2-}\), \(Cr_2O_7^{2-}\), \(MnO_4^{2-}\), \(ZnO_2^{2-}\), \(MoO_4^{2-}\), \(SnO_2^{2-}\), \(SnO_3^{2-}\), \(SiO_3^{2-}\)
-3
\(B^{3-}\), \(N^{3-}\), \(BO_3^{3-}\), \(PO_3^{3-}\), \(PO_4^{3-}\), \(AlO_3^{3-}\), \(As^{3-}\), \(AsO_3^{3-}\), \(AsO_4^{3-}\), \(SbO_4^{3-}\)
-4
\(C^{4-}\), \(P_2O_4^{4-}\), \([Fe(CN)_6]^{4-}\)
❖ Halogen Oxyanion Series:

Table 1: Halogen Oxyanions

Halogen
Series
Chlorine
\(Cl^-\), \(ClO^-\), \(ClO_2^-\), \(ClO_3^-\), \(ClO_4^-\)
Bromine
\(Br^-\), \(BrO^-\), \(BrO_2^-\), \(BrO_3^-\), \(BrO_4^-\)
Iodine
\(I^-\), \(IO^-\), \(IO_2^-\), \(IO_3^-\), \(IO_4^-\)
❖ Iron Cyanide Radicals:

Table 1: Complex Cyanide Radicals

Name
Formula
Ferrocyanide
\([Fe(CN)_6]^{4-}\)
Ferricyanide
\([Fe(CN)_6]^{3-}\)
▢ Catalyst:
❖ Definition: Substance altering rate of chemical reaction without change in mass and composition
❖ Phenomenon: Catalysis
❖ Types:

Table 1: Types of Catalyst

Type
Effect
Activation energy
Example
Positive catalyst
Increases reaction rate
Decreases
\(MnO_2\) in decomposition of \(KClO_3\)
Negative catalyst
Decreases reaction rate
Increases
Dil. \(H_2SO_4\) / glycerol in \(H_2O_2\) decomposition
Autocatalyst
Product itself acts as catalyst
Self-catalysis
\(Mn^{2+}\) in \(KMnO_4\) vs oxalic acid titration
❖ Positive Catalyst Example:
  • •
    Without catalyst: \(2KClO_3 \xrightarrow{400^\circ C} 2KCl + 3O_2\) → slow
  • •
    With catalyst: \(2KClO_3 \xrightarrow[MnO_2]{250^\circ C} 2KCl + 3O_2\) → fast
❖ Negative Catalyst Example:
  • •
    \(2H_2O_2 \rightarrow 2H_2O + O_2\) → fast
  • •
    \(2H_2O_2 \xrightarrow{Pt} 2H_2O + O_2\) → very fast
  • •
    \(2H_2O_2 \xrightarrow{Dil.\ H_2SO_4 / Glycerol} 2H_2O + O_2\) → slow
❖ Autocatalysis Examples:
  • •
    \(2MnO_4^- + 5C_2O_4^{2-} + 16H^+ \rightarrow 2Mn^{2+} + 10CO_2 + 8H_2O\) → \(Mn^{2+}\) acts as catalyst
  • •
    \(CH_3COOC_2H_5 + H_2O \rightarrow CH_3COOH + C_2H_5OH\) → acetic acid acts as autocatalyst
❖ Characteristics:
  • •
    Only alters rate of reaction
  • •
    Mass + composition unchanged at end
  • •
    Small amount sufficient
  • •
    Cannot initiate reaction
  • •
    Specific in action
❖ Notes:
  • •
    Acid–base reaction = double displacement reaction
  • •
    Symbol = abbreviation of element name
  • •
    Sand + powdered charcoal separated by washing
▢ Atomic Weight / Atomic Mass:
❖ Historical Reference:
  • •
    First reference → hydrogen / Dalton
  • •
    Later reference → oxygen-16
  • •
    Present reference → carbon-12
❖ Definition: Number showing how many times heavier an average atom of element is than \(\frac{1}{12}\) mass of \(^{12}C\)
❖ Formula: \(Atomic\ weight = \frac{Average\ weight\ of\ atom}{\frac{1}{12} \times mass\ of\ ^{12}C\ atom}\)
❖ Nature: Average relative weight
▢ Unit of Atomic Weight:

Table 1: Atomic Mass Unit

Point
Value
Modern unit
\(\frac{1}{12}\) mass of \(^{12}C\) atom
Atomic weight
Ratio; no true unit but expressed in amu
1 amu
\(\frac{1}{12}\times mass\ of\ ^{12}C\ atom\)
Mass of one \(^{12}C\) atom
\(1.9924\times10^{-23}\ g\)
1 amu in gram
\(1.66\times10^{-24}\ g\)
Lightest atom
Hydrogen
Heaviest naturally occurring atom
\(^{238}U\)
❖ Fractional Atomic Weight:
  • •
    Due to stable isotopes
  • •
    Depends on mass numbers of isotopes
  • •
    Depends on relative abundance of isotopes
❖ Chlorine Example: \(\frac{75\times35 + 25\times37}{100} = 35.5\ amu\)
▢ Gram Atomic Weight:
❖ Definition: Atomic weight expressed in grams
❖ Example: Atomic weight of sulphur = 32; gram atomic weight = 32 g
▢ Gram Atom / Mole Atom:
❖ Definition: Atomic weight of element in grams = 1 gram atom / 1 mole atom
❖ Example: 1 gram atom of sodium = 23 g sodium
❖ Composition Terms:

Table 1: Composition of Water

Type
Meaning
Example
Gravimetric composition
Composition by weight
Water → H:O = 1:8
Volumetric composition
Composition by volume
Water → H:O = 2:1
▢ Molecular Weight:
❖ Definition: Sum of atomic weights of all atoms present in one molecule
❖ Example: \(H_2O = 2\times1 + 16 = 18\)
▢ Equivalent Weight:
❖ Definition: Combining weight of a substance relative to fixed amount of hydrogen / oxygen / chlorine
❖ General Formula: \(Equivalent\ weight = \frac{Molecular\ weight}{n-factor}\)
❖ For Acids: \(Equivalent\ weight = \frac{Molecular\ weight}{basicity}\)
❖ For Bases: \(Equivalent\ weight = \frac{Molecular\ weight}{acidity}\)
❖ For Salts: \(Equivalent\ weight = \frac{Molecular\ weight}{total\ positive\ valency\ or\ total\ negative\ valency}\)
❖ Note: Detailed table after this point was incomplete in uploaded content
▢ High-Yield Recall:

Table 1: Basic Chemistry One-Liners

Fact
Answer
Matter–energy equation
\(E=mc^2\)
Element concept
Robert Boyle
Atom term
Dalton
Molecule term
Avogadro
Compound
Elements in fixed proportion
Mixture
Substances in any proportion
Compound separation
Chemical method
Mixture separation
Physical method
Atmolysis
Separation of gases / vapours by diffusion rate
Substrate
Substance attacked by reagent
Reagent
Attacking species
Precipitate
Insoluble product from solutions
Balanced equation based on
Conservation of mass
Positive radical
Cation / basic radical
Negative radical
Anion / acid radical
Cation moves toward
Cathode
Basic radical source
Base
Acid radical source
Acid
Positive catalyst
Increases reaction rate; decreases activation energy
Negative catalyst
Decreases reaction rate; increases activation energy
Autocatalyst example
\(Mn^{2+}\) in acidified \(KMnO_4\) + oxalic acid reaction
Present atomic mass standard
\(^{12}C\)
1 amu
\(1.66\times10^{-24}\ g\)
Lightest atom
Hydrogen
Heaviest natural atom
\(^{238}U\)
Chlorine atomic weight
35.5 amu
Gram atomic weight
Atomic weight in grams
1 gram atom Na
23 g sodium
Water gravimetric ratio
H:O = 1:8
Water volumetric ratio
H:O = 2:1
Equivalent weight formula
\(Molecular\ weight / n-factor\)
Q1.
Which one of the following pairs of compounds illustrate the law of multiple proportions?
Q2.
The hydrogen phosphate of certain metal has formula MHPO4. The formula of metal chloride would be
Q3.
Modern atomic weight scale is based on
Q4.
1 amu is equal to
Q5.
The ratio of the rates of diffusion of a given element to that of helium is 4. What will be the molecular weight of the element?
Q6.
The law of reciprocal proportions can be used to determine
Q7.
The specific heat of a metal is 0.16, its approximate atomic weight would be
Q8.
If the equivalent weight of a trivalent metal is 32.7, the molecular weight of its chloride is
Q9.
Vapour density of a gas is 22. Its molecular weight will be
Q10.
In the reaction: 2Na2S2O3 + I2 → Na2S4O6 + 2NaI, the equivalent weight of Na2S2O3 (mol. wt. = M) is equal to
Q11.
A metallic oxide contains 60% of the metal. The equivalent weight of the metal is
Q12.
The equivalent weight of KMnO4 (in acid medium) is (At. wt. of K = 39, Mn = 55)
Q13.
M is the molecular weight of KMnO4. The equivalent weight of KMnO4 when it is converted into K2MnO4 is
Q14.
In the reaction: SO2 + 2H2S → 3S + 2H2O, which choice has value twice that of the equivalent weight of the oxidising agent?
Q15.
Equivalent weight of an acid whose basicity is 3, is
Q16.
The number of water molecules in 1 litre of water is
Q17.
The number of water molecules present in a drop of water (volume = 0.0018 ml) at room temperature is
Q18.
Which has maximum molecules?
Q19.
Volume of 4.4 g of CO2 at STP is
Q20.
The empirical formula of an acid is CH2O2, the probable molecular formula of the acid may be
Q21.
The percentage of nitrogen in urea (NH2)2CO is about
Q22.
1.25 g NH3 contains how many atoms?
Q23.
1 gm atom of nitrogen contains
Q24.
If isotopic distribution of C-12 and C-14 is 98% and 2% respectively, then the number of C-14 atoms in 12 g of carbon is:
Q25.
Which one of the following gases contains the same number of molecules as 16g of oxygen?
Q26.
Atomic weights of some elements are not whole numbers because
Q27.
Which property of element is always a whole number?
Q28.
Which pair among the following crystals are isomorphous?
Q29.
Law of definite proportion does not apply to nitrogen oxide because
Q30.
The simplest formula of a compound containing 50% of element X (atomic wt. 10) and 50% of element Y (atomic wt. 20) is