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By rohit.pandey1
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Updated on 8 Jul 2026, 14:43 IST
A Class 11 Chemistry formula sheet PDF helps students revise important formulas quickly before school exams, unit tests, JEE foundation practice and NEET foundation preparation. This chapter-wise guide covers the most important formulas from Some Basic Concepts of Chemistry, Structure of Atom, Classification of Elements, Chemical Bonding, Thermodynamics, Equilibrium, Redox Reactions, Organic Chemistry and Hydrocarbons.
This article is written for students who want a clean, accurate and easy-to-revise collection of chemistry formulas for Class 11. Each formula is given with its meaning, condition, unit or important exception wherever required.
CBSE syllabus note: CBSE Class XI Chemistry is arranged in units in the official course structure. The CBSE Class XI Chemistry theory paper is listed for 70 marks and includes units such as Some Basic Concepts of Chemistry, Structure of Atom, Classification of Elements and Periodicity in Properties, Chemical Bonding and Molecular Structure, Chemical Thermodynamics, Equilibrium, Redox Reactions, Organic Chemistry and Hydrocarbons. These values are CBSE unit marks/course structure, not guaranteed chapter-wise weightage in the final question paper.
The most important formulas in Class 11 Chemistry come from mole concept, concentration terms, atomic structure, chemical bonding, thermodynamics, equilibrium, redox reactions, organic analysis and hydrocarbons. Use this formula sheet for revision, but always learn the conditions, units and exceptions along with each formula.
The best way to study Class 11 Chemistry formulas is to revise them chapter wise. This helps students connect formulas with the exact chapter and concept instead of memorising random equations.
| Sr. No. | Chapter Name | Download PDF |
| 1 | Some Basic Concepts of Chemistry | Download PDF |
| 2 | Structure of Atom | Download PDF |
| 3 | Classification of Elements & Periodicity in Properties | Download PDF |
| 4 | Some Basic Concepts of Chemistry | Download PDF |
| 5 | States of Matter | Download PDF |
| 6 | Thermodynamics | Download PDF |
| 7 | Equilibrium | Download PDF |
| 8 | Redox Reactions | Download PDF |
| 9 | Hydrogen | Download PDF |
| 10 | The s-Block Elements | Download PDF |
| 11 | The p-Block Elements | Download PDF |
| 12 | Organic Chemistry – Some Basic Principles and Techniques | Download PDF |
Download Class 11 Chemistry Important Formulas Chapter Wise PDF to revise all key formulas in one place. This PDF includes important formulas from chapters like Some Basic Concepts of Chemistry, Structure of Atom, Chemical Bonding, Thermodynamics, Equilibrium, Redox Reactions, Organic Chemistry, and Hydrocarbons. Students can use it for quick revision before school exams, CBSE tests, JEE foundation, and NEET preparation. Each formula is arranged chapter-wise to make learning and practice easier.
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Some Basic Concepts of Chemistry is the foundation of Class 11 Physical Chemistry. Mole concept, stoichiometry and concentration terms are used repeatedly in later chapters such as Thermodynamics, Equilibrium, Redox Reactions and Organic Chemistry.
| Formula | Meaning | Unit / Note |
| n = m / M | Number of moles | n = moles, m = given mass, M = molar mass |
| m = nM | Mass from moles | Usually in grams |
| Number of particles = nNA | Atoms, molecules or ions from moles | NA = 6.022 × 1023 mol−1 |
| n = Number of particles / NA | Moles from particles | mol |
| M = m / n | Molar mass | g mol−1 |
The most important formulas of Chemistry Class 11 Chapter 1 are mole formula, Avogadro relation, molarity, molality, mole fraction, percentage composition, empirical formula, molecular formula, limiting reagent and percentage yield.
| Formula | Use |
| % of element = (mass of element in compound / molar mass of compound) × 100 | Finds percentage of an element in a compound |
| Mass of element = (% of element × molar mass of compound) / 100 | Finds mass contribution of an element |
| Formula | Meaning |
| Empirical formula mass = sum of atomic masses in empirical formula | Mass of simplest formula |
| Multiplying factor = molecular mass / empirical formula mass | Whole-number multiplier |
| Molecular formula = empirical formula × multiplying factor | Actual molecular formula |
Exam tip: In empirical formula questions, convert percentage or mass into moles, divide each mole value by the smallest mole value and convert the ratio into whole numbers.
| Term | Formula | Unit |
| Molarity | M = moles of solute / volume of solution in litre | mol L−1; temperature-dependent |
| Molality | m = moles of solute / mass of solvent in kg | mol kg−1; temperature-independent |
| Normality | N = gram equivalents of solute / volume of solution in litre | eq L−1; reaction-dependent |
| Relation between normality and molarity | N = M × n-factor | n-factor depends on reaction |
| Mole fraction | XA = nA / Σn | Unitless |
| Mass percentage | % w/w = mass of solute / mass of solution × 100 | % |
| Volume percentage | % v/v = volume of solute / volume of solution × 100 | % |
| Parts per million | ppm = mass of solute / mass of solution × 106 | For very dilute solutions |
| Dilution | M1V1 = M2V2 | Same solute before and after dilution |
| Formula / Method | Use |
| Mole ratio = coefficient ratio from balanced equation | Compares reactants and products |
| Theoretical yield = product amount calculated from balanced equation | Maximum possible yield |
| % yield = actual yield / theoretical yield × 100 | Reaction efficiency |
| % purity = mass of pure substance / mass of impure sample × 100 | Purity calculation |
| Limiting reagent = reactant that produces the least product | Reactant consumed first |
| Mistake | Correct Approach |
| Using mass directly in stoichiometry | Convert mass into moles first |
| Not balancing the equation | Balance before using mole ratio |
| Confusing molarity and molality | Molarity uses solution volume; molality uses solvent mass |
| Using grams instead of kg in molality | Convert solvent mass to kg |
| Treating normality as fixed | Normality depends on n-factor and reaction type |
Structure of Atom includes electromagnetic radiation, Bohr’s model, hydrogen-like ions, de Broglie wavelength, Heisenberg uncertainty principle and quantum numbers.
| Formula | Meaning |
| Z = number of protons | Atomic number |
| A = number of protons + number of neutrons | Mass number |
| Number of neutrons = A − Z | Neutron count |
| For a neutral atom: protons = electrons | Charge balance |
| Formula | Meaning |
| E = hν | Energy of one photon |
| c = νλ | Relation between speed, frequency and wavelength |
| E = hc / λ | Photon energy in terms of wavelength |
Here, E is energy, h is Planck’s constant, ν is frequency, λ is wavelength and c is the speed of light.
Exam tip: Frequency and wavelength are inversely related. If wavelength increases, frequency decreases.
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For hydrogen-like species such as H, He+, Li2+ and Be3+, use the following formulas.
| Formula | Meaning |
| En = −13.6Z2 / n2 eV | Energy of electron in nth orbit |
| rn = 0.529n2 / Z Å | Radius of nth orbit |
| mvr = nh / 2π | Bohr quantisation of angular momentum |
For hydrogen atom, Z = 1. Therefore:
| Hydrogen Atom Formula | Meaning |
| En = −13.6 / n2 eV | Energy of electron in nth orbit of hydrogen |
| rn = 0.529n2 Å | Radius of nth orbit of hydrogen |
For an electronic transition, use:
hν = |Ef − Ei|
Here, Ei is the initial energy level, Ef is the final energy level and hν is the photon energy.
In emission, the electron moves from a higher energy level to a lower energy level. Therefore, Ef < Ei, and energy is released as radiation.
In absorption, the electron moves from a lower energy level to a higher energy level. Therefore, Ef > Ei, and energy is absorbed.
Important: Photon energy is always positive, so use the magnitude: hν = |Ef − Ei|.
For hydrogen-like species:
1 / λ = RZ2(1 / n12 − 1 / n22)
For emission, n2 > n1, where the electron falls from n2 to n1.
| Formula | Meaning |
| λ = h / p | Wavelength from momentum |
| λ = h / mv | Wavelength of a moving particle |
| λ = h / √(2mEk) | Wavelength from kinetic energy |
| Formula | Meaning |
| Δx · Δp ≥ h / 4π | Position-momentum uncertainty |
| Δx · mΔv ≥ h / 4π | Position-velocity uncertainty |
| Δx · Δp ≥ ℏ / 2 | Equivalent form, where ℏ = h / 2π |
| Quantum Number | Symbol | Allowed Values |
| Principal quantum number | n | 1, 2, 3, ... |
| Azimuthal quantum number | l | 0 to n − 1 |
| Magnetic quantum number | ml | −l to +l |
| Spin quantum number | ms | +1/2 or −1/2 |
| Subshell | l Value |
| s | 0 |
| p | 1 |
| d | 2 |
| f | 3 |
This unit is not formula-heavy, but it is important for reasoning-based questions. Students should focus on periodic trends, exceptions and the meaning of effective nuclear charge.
| Property | Across a Period | Down a Group |
| Atomic radius | Generally decreases | Increases |
| Ionic radius | Generally decreases for similar charge type | Increases |
| Ionization enthalpy | Generally increases | Generally decreases |
| Electron gain enthalpy | Generally becomes more negative | Generally becomes less negative |
| Electronegativity | Generally increases | Generally decreases |
| Metallic character | Decreases | Increases |
| Non-metallic character | Increases | Decreases |
A simplified relation is:
Zeff = Z − σ
Here, Zeff is effective nuclear charge, Z is atomic number and σ is shielding or screening constant.
Note: This is a simplified trend-based relation. Exact calculations may require more advanced methods.
Chemical Bonding formulas help in Lewis structures, formal charge, molecular orbital theory, dipole moment and hybridisation.
Formal charge = V − N − B/2
Where V is valence electrons of the free atom, N is non-bonding electrons and B is bonding electrons.
A stable Lewis structure usually has minimum formal charges, and negative charge is more stable on the more electronegative atom.
Bond order = (Nb − Na) / 2
Where Nb is the number of electrons in bonding molecular orbitals and Na is the number of electrons in antibonding molecular orbitals.
| Bond Order | Meaning |
| 0 | Molecule is usually unstable |
| 1 | Single bond |
| 2 | Double bond |
| 3 | Triple bond |
μ = q × r
Where μ is dipole moment, q is charge and r is distance between charges.
Unit conversion: 1 D = 3.336 × 10−30 C m
Exam tip: A molecule can have polar bonds but zero net dipole moment if the molecular geometry is symmetrical.
Steric number = number of σ bonds + number of lone pairs
| Steric Number | Hybridisation | Common Shape |
| 2 | sp | Linear |
| 3 | sp2 | Trigonal planar |
| 4 | sp3 | Tetrahedral |
| 5 | sp3d | Trigonal bipyramidal |
| 6 | sp3d2 | Octahedral |
| Mistake | Correction |
| Counting π bonds in steric number | Count only σ bonds and lone pairs |
| Ignoring antibonding electrons | Use both bonding and antibonding electrons in bond order |
| Assuming every polar bond makes a polar molecule | Check molecular shape |
| Forgetting formal charge | Use formal charge to compare Lewis structures |
Thermodynamics connects heat, work, internal energy, enthalpy, entropy, Gibbs energy and spontaneity. This chapter is highly important for Class 11 Chemistry numericals.
ΔU = q + w
This formula uses the chemistry sign convention.
| Quantity | Meaning |
| q > 0 | Heat absorbed by the system |
| q < 0 | Heat released by the system |
| w > 0 | Work done on the system |
| w < 0 | Work done by the system |
| Formula | Condition |
| w = −PextΔV | Only for constant external pressure |
| ΔV = V2 − V1 | Change in volume |
| wrev = −nRT ln(V2/V1) | Reversible isothermal expansion or compression of an ideal gas |
| wrev = −2.303nRT log(V2/V1) | Same formula using common logarithm |
Exam tip: During expansion, V2 > V1, so work is negative under chemistry convention.
| Formula | Meaning / Condition |
| H = U + PV | Definition of enthalpy |
| ΔH = ΔU + Δ(PV) | General relation |
| ΔH = ΔU + ΔngRT | For ideal gaseous reactions |
| qp = ΔH | Heat at constant pressure |
| qv = ΔU | Heat at constant volume |
Here, Δng = moles of gaseous products − moles of gaseous reactants. Count only gaseous species.
| Formula | Meaning |
| q = mcΔT | Heat change using specific heat |
| C = q / ΔT | Heat capacity |
| c = q / mΔT | Specific heat capacity |
| Cm = q / nΔT | Molar heat capacity |
ΔrH° = ΣνΔfH°(products) − ΣνΔfH°(reactants)
Here, ΔrH° is standard enthalpy change of reaction, ΔfH° is standard enthalpy of formation and ν is the stoichiometric coefficient.
General definition:
ΔS = ∫δqrev / T
For a reversible isothermal process only:
ΔS = qrev / T
Total entropy change:
ΔSuniverse = ΔSsystem + ΔSsurroundings
Important: Do not use ΔS = q/T blindly. The expression ΔS = qrev/T applies only when the process is reversible and isothermal.
| Formula | Meaning |
| ΔG = ΔH − TΔS | Gibbs energy change |
| ΔG < 0 | Spontaneous process |
| ΔG > 0 | Non-spontaneous process |
| ΔG = 0 | Equilibrium |
| ΔG = ΔG° + RT ln Q | Gibbs energy under non-standard conditions |
| ΔG° = −RT ln K | Relation between standard Gibbs energy and equilibrium constant |
Important conditions: T must be in Kelvin. K should be treated as dimensionless in ΔG° = −RT ln K. At equilibrium, ΔG = 0 and Q = K.
| ΔH | ΔS | Spontaneity |
| Negative | Positive | Spontaneous at all temperatures |
| Positive | Negative | Non-spontaneous at all temperatures |
| Negative | Negative | Spontaneous at low temperature |
| Positive | Positive | Spontaneous at high temperature |
| Mistake | Correct Approach |
| Using Celsius in RT or TΔS | Use Kelvin |
| Applying w = −PextΔV to all processes | Use only for constant external pressure |
| Writing ΔS = q/T without condition | Use ΔS = ∫δqrev/T generally |
| Ignoring chemistry sign convention | Work done by the system is negative |
| Using ΔG° = −RT ln K without conditions | K is dimensionless and T is in Kelvin |
Equilibrium includes chemical equilibrium, ionic equilibrium, acid-base equilibrium, buffer solutions and solubility product. This chapter is one of the most important formula-based chapters in Class 11 Chemistry.
For the reaction:
aA + bB ⇌ cC + dD
The concentration equilibrium constant is:
Kc = [C]c[D]d / [A]a[B]b
The pressure equilibrium constant is:
Kp = (PC)c(PD)d / (PA)a(PB)b
Important conditions:
The reaction quotient Q has the same form as K, but it uses concentrations or pressures at any stage of the reaction.
| Condition | Direction |
| Q < K | Reaction proceeds forward |
| Q > K | Reaction proceeds backward |
| Q = K | System is at equilibrium |
Kp = Kc(RT)Δng
Where:
Δng = moles of gaseous products − moles of gaseous reactants
Condition: This relation is valid for ideal gases.
Exam tip: Count only gaseous species while calculating Δng. Do not count solids, liquids or aqueous species.
| Formula | Meaning |
| Ka = [H+][A−] / [HA] | Acid dissociation constant |
| Kb = [BH+][OH−] / [B] | Base dissociation constant |
| Kw = [H+][OH−] | Ionic product of water |
| KaKb = Kw | For a conjugate acid-base pair |
| pKa = −log Ka | Acid strength scale |
| pKb = −log Kb | Base strength scale |
| pKw = −log Kw | Water ionisation scale |
| Formula | Meaning |
| pH = −log[H+] | Hydrogen ion concentration scale |
| pOH = −log[OH−] | Hydroxide ion concentration scale |
| pH + pOH = pKw | General relation |
| At 25°C, pKw = 14 | Therefore pH + pOH = 14 at 25°C |
| [H+] = 10−pH | Hydrogen ion concentration from pH |
| [OH−] = 10−pOH | Hydroxide ion concentration from pOH |
Important: pH + pOH = 14 is true at 25°C. The general relation is pH + pOH = pKw.
For an acidic buffer:
pH = pKa + log([salt] / [acid])
More specifically:
pH = pKa + log([A−] / [HA])
For a basic buffer:
pOH = pKb + log([salt] / [base])
More specifically:
pOH = pKb + log([BH+] / [B])
Validity conditions:
For sparingly soluble salts:
Ksp = product of ion concentrations raised to their stoichiometric powers
| Salt Dissociation | Ksp Relation |
| AB ⇌ A+ + B− | Ksp = s2 |
| AB2 ⇌ A2+ + 2B− | Ksp = s(2s)2 = 4s3 |
| A2B ⇌ 2A+ + B2− | Ksp = (2s)2s = 4s3 |
| AB3 ⇌ A3+ + 3B− | Ksp = s(3s)3 = 27s4 |
These simplified relations assume pure water, molar solubility s, no common ion and complete dissociation.
| Mistake | Correction |
| Using initial concentration in Kc | Use equilibrium concentration |
| Including pure solids and liquids | Omit pure solids and pure liquids |
| Writing pH + pOH = 14 without condition | Write pH + pOH = pKw; at 25°C, pKw = 14 |
| Applying Kp = Kc(RT)Δng to non-gaseous systems | Use it for ideal gases |
| Forgetting powers in K expression | Use balanced equation coefficients as powers |
Redox formulas are used in oxidation number, equivalent weight, n-factor, normality and balancing redox equations.
| Rule | Example / Exception |
| Free element has oxidation number 0 | Na, O2, Cl2 |
| Monoatomic ion has oxidation number equal to charge | Na+ = +1, Cl− = −1 |
| Oxygen is usually −2 | Common oxides |
| Oxygen in peroxides is −1 | H2O2, Na2O2 |
| Oxygen in superoxides is −1/2 | KO2 |
| Oxygen in OF2 is +2 | Fluorine is more electronegative |
| Oxygen in O2F2 is +1 | Oxygen fluoride exception |
| Hydrogen is usually +1 | Except metal hydrides, where H is −1 |
| Sum of oxidation numbers in a neutral compound is 0 | H2O, CO2 |
| Sum in a polyatomic ion equals ion charge | SO42−, NO3− |
| Formula | Meaning |
| Equivalent weight = molar mass / n-factor | Equivalent mass |
| Number of equivalents = given mass / equivalent weight | Equivalent count |
| N = equivalents / volume in litre | Normality |
| N = M × n-factor | Relation between normality and molarity |
| Milliequivalents = N × VmL | Useful in titration-style problems |
| Species Type | n-Factor Meaning |
| Acid | Number of replaceable H+ ions in the given reaction |
| Base | Number of replaceable OH− ions or acid-neutralising capacity in the given reaction |
| Salt | Total charge exchanged, depending on reaction |
| Redox species | Change in oxidation number per formula unit in the given reaction |
Important: n-factor is reaction-dependent, especially for redox species. The same compound may have different n-factors in different reactions.
Use either the oxidation-number method or the half-reaction method.
Organic Chemistry in Class 11 includes structure, nomenclature, isomerism, reaction intermediates, purification and organic analysis.
For a compound containing carbon, hydrogen, nitrogen and halogens:
DBE = C − (H + X)/2 + N/2 + 1
Where C is the number of carbon atoms, H is the number of hydrogen atoms, X is the number of halogen atoms and N is the number of nitrogen atoms. Oxygen and sulfur are ignored in this formula.
Exam tip: Halogens are counted with hydrogen in DBE. Use H + X, not just H.
| Formula | Use |
| % of element = mass of element / total mass of compound × 100 | Elemental composition |
| Empirical formula mass = sum of atomic masses in empirical formula | Simplest formula mass |
| Multiplying factor = molecular mass / empirical formula mass | Molecular formula factor |
| Molecular formula = empirical formula × multiplying factor | Actual molecular formula |
| Term | Correct Meaning |
| Electrophile | Electron-deficient species or electron-pair acceptor |
| Nucleophile | Electron-rich species/electron-pair donor that attacks electron-deficient centres |
| Homolytic fission | Covalent bond breaks equally, forming free radicals |
| Heterolytic fission | Covalent bond breaks unequally, forming ions |
| Inductive effect | Electron displacement through σ bonds |
| Resonance effect | Delocalisation of π electrons or lone pairs |
| Hyperconjugation | Delocalisation involving σ electrons of C-H or C-C bonds adjacent to an empty or partially filled p-orbital or π system |
| Element | Formula |
| Carbon | %C = (12 × mass of CO2 × 100) / (44 × mass of organic compound) |
| Hydrogen | %H = (2 × mass of H2O × 100) / (18 × mass of organic compound) |
If nitrogen gas volume is measured at STP in mL:
%N = (28 × VN2 × 100) / (22400 × w)
Here, VN2 is the volume of nitrogen gas at STP in mL and w is the mass of organic compound in grams.
%N = 1.4NV / w
Here, N is normality of acid, V is volume of acid used in mL and w is mass of organic compound in grams.
Limitation: Kjeldahl method is not suitable for all nitrogen-containing compounds, especially compounds where nitrogen is present in nitro or azo groups.
%X = (atomic mass of X × mass of AgX × 100) / (molar mass of AgX × mass of organic compound)
Here, X may be Cl, Br or I.
%S = (32 × mass of BaSO4 × 100) / (233 × mass of organic compound)
%P = (62 × mass of Mg2P2O7 × 100) / (222 × mass of organic compound)
Note: In numerical problems, use the atomic masses and molar masses given in the question or textbook table if they differ slightly from rounded values.
Hydrocarbons contain only carbon and hydrogen. Formula-based questions usually involve general formulas, unsaturation, combustion and reaction patterns.
| Series | General Formula | Condition / Example |
| Alkane | CnH2n+2 | Acyclic saturated hydrocarbon |
| Alkene | CnH2n | Acyclic hydrocarbon with one double bond |
| Alkyne | CnH2n−2 | Acyclic hydrocarbon with one triple bond |
| Cycloalkane | CnH2n | Saturated monocyclic hydrocarbon |
| Monocyclic arene / benzene homologue | CnH2n−6 | Benzene, toluene series |
Important correction: CnH2n−6 applies to monocyclic arenes or benzene homologues. It should not be used for every aromatic hydrocarbon.
For complete combustion:
CxHy + (x + y/4)O2 → xCO2 + (y/2)H2O
This formula is useful for quickly balancing complete combustion reactions of hydrocarbons.
| Reaction | General Pattern |
| Hydrogenation | Alkene + H2 → alkane |
| Halogenation | Alkene + X2 → dihaloalkane |
| Hydrohalogenation | Alkene + HX → alkyl halide |
| Hydration | Alkene + H2O → alcohol |
| Complete combustion | Hydrocarbon + O2 → CO2 + H2O |
| Topic | Formula |
| Mole concept | n = m / M |
| Avogadro relation | Number of particles = nNA |
| Molarity | M = moles of solute / volume of solution in litre |
| Molality | m = moles of solute / mass of solvent in kg |
| Photon energy | E = hν = hc / λ |
| Wave relation | c = νλ |
| Bohr energy | En = −13.6Z2 / n2 eV |
| Bohr radius | rn = 0.529n2 / Z Å |
| de Broglie wavelength | λ = h / mv |
| Uncertainty principle | ΔxΔp ≥ h / 4π |
| Formal charge | V − N − B/2 |
| Bond order | (Nb − Na) / 2 |
| First law | ΔU = q + w |
| Constant-pressure work | w = −PextΔV |
| Reversible isothermal work | wrev = −nRT ln(V2/V1) |
| Entropy | ΔS = ∫δqrev / T |
| Gibbs energy | ΔG = ΔH − TΔS |
| Non-standard Gibbs energy | ΔG = ΔG° + RT ln Q |
| Equilibrium Gibbs relation | ΔG° = −RT ln K |
| Kp-Kc relation | Kp = Kc(RT)Δng |
| pH | pH = −log[H+] |
| pOH | pOH = −log[OH−] |
| pH-pOH relation | pH + pOH = pKw |
| Equivalent weight | molar mass / n-factor |
| DBE | C − (H + X)/2 + N/2 + 1 |
| Alkane | CnH2n+2 |
| Alkene | CnH2n |
| Alkyne | CnH2n−2 |
Use this formula sheet in three rounds. First, understand what every symbol means. Second, write the formula with units. Third, solve one example based on the formula. This method is better than memorising equations without knowing where they are used.
Exam tip: Most students lose marks not because they forget the formula, but because they use the wrong unit, wrong sign convention, wrong concentration term or wrong condition.
The table below shows the CBSE unit marks/course structure. Do not call these values chapter-wise weightage because the actual question paper may not distribute questions exactly according to this table.
| Unit | CBSE Unit / Topic | Formula Focus |
| 1 | Some Basic Concepts of Chemistry | Mole concept, stoichiometry, concentration terms |
| 2 | Structure of Atom | Radiation, Bohr model, de Broglie relation, uncertainty principle |
| 3 | Classification of Elements and Periodicity in Properties | Periodic trends, effective nuclear charge idea |
| 4 | Chemical Bonding and Molecular Structure | Formal charge, bond order, dipole moment, hybridisation |
| 5 | Chemical Thermodynamics | Internal energy, work, enthalpy, entropy, Gibbs energy |
| 6 | Equilibrium | Kc, Kp, Q, pH, pOH, buffer, Ksp |
| 7 | Redox Reactions | Oxidation number, n-factor, equivalent weight |
| 8 | Organic Chemistry: Some Basic Principles and Techniques | DBE, empirical formula, molecular formula, organic analysis |
| 9 | Hydrocarbons | General formulas, combustion, reaction patterns |
| Mistake | Correct Approach |
| Using mL instead of litre in molarity | Convert mL to L |
| Using gram instead of kg in molality | Convert solvent mass to kg |
| Using Celsius in thermodynamic formulas | Use Kelvin |
| Forgetting chemistry sign convention | In expansion, work done by system is negative |
| Using pH + pOH = 14 universally | Use pH + pOH = pKw; at 25°C, pKw = 14 |
| Including solids in equilibrium constants | Omit pure solids and pure liquids |
| Applying Kp = Kc(RT)Δng to all systems | Use it for ideal gases |
| Forgetting halogens in DBE | Use H + X |
| Treating n-factor as fixed | n-factor depends on the reaction |
| Using CnH2n−6 for all aromatics | Use it only for monocyclic arenes / benzene homologues |
This Class 11 Chemistry important formula sheet PDF-style guide gives the key formulas needed for CBSE Class 11 Chemistry, NCERT revision, JEE foundation and NEET foundation preparation. The best way to use it is chapter-wise: first revise the formula, then understand the condition, then solve one question based on it.
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The most important formulas in Class 11 Chemistry include mole concept formulas, concentration terms, atomic structure formulas, formal charge, bond order, thermodynamics formulas, equilibrium formulas, redox n-factor formulas, DBE and hydrocarbon general formulas.
Important Chapter 1 formulas include n = m/M, particles = nNA, molarity, molality, mole fraction, percentage composition, empirical formula, molecular formula, limiting reagent and percentage yield.
The formula for number of moles is n = m/M, where n is number of moles, m is given mass and M is molar mass.
The formula for molarity is M = moles of solute / volume of solution in litre. Molarity is expressed in mol L−1.
The formula for molality is m = moles of solute / mass of solvent in kg. Molality is expressed in mol kg−1.
Important formulas in Structure of Atom include E = hν, c = νλ, E = hc/λ, En = −13.6Z2/n2 eV, rn = 0.529n2/Z Å, λ = h/mv and ΔxΔp ≥ h/4π.
de Broglie’s equation is λ = h/mv. It relates the wavelength of a moving particle to its mass and velocity.
Heisenberg’s uncertainty principle is ΔxΔp ≥ h/4π. It states that the exact position and exact momentum of a microscopic particle cannot be known simultaneously.
Important thermodynamics formulas include ΔU = q + w, w = −PextΔV, wrev = −nRT ln(V2/V1), ΔH = ΔU + ΔngRT, ΔS = ∫δqrev/T, ΔG = ΔH − TΔS and ΔG = ΔG° + RT ln Q.
The first law of thermodynamics is ΔU = q + w. It uses the chemistry sign convention, where heat absorbed by the system is positive and work done on the system is positive.
For constant external pressure, work is w = −PextΔV. For reversible isothermal expansion or compression of an ideal gas, work is wrev = −nRT ln(V2/V1).
Important equilibrium formulas include Kc, Kp, Kp = Kc(RT)Δng, pH = −log[H+], pOH = −log[OH−], pH + pOH = pKw, Henderson–Hasselbalch equation and Ksp.
The relation is Kp = Kc(RT)Δng. This relation applies to ideal gases. Here, Δng is moles of gaseous products minus moles of gaseous reactants.
No. The general relation is pH + pOH = pKw. At 25°C, pKw = 14, so pH + pOH = 14 at 25°C.
For an acidic buffer, the Henderson–Hasselbalch equation is pH = pKa + log([A−]/[HA]). For a basic buffer, pOH = pKb + log([BH+]/[B]). It is valid for buffer solutions where both conjugate acid-base components are present in appreciable amounts.
Important Redox formulas include equivalent weight = molar mass / n-factor, N = M × n-factor and milliequivalents = N × VmL. The n-factor is reaction-dependent, especially in redox reactions.
The DBE formula is DBE = C − (H + X)/2 + N/2 + 1. Here X represents halogens. Oxygen and sulfur are ignored in this formula.
The general formula of alkanes is CnH2n+2, the general formula of alkenes is CnH2n and the general formula of alkynes is CnH2n−2.
No. CnH2n−6 applies to monocyclic arenes or benzene homologues, not every aromatic hydrocarbon.
No. Formula revision is useful, but students also need NCERT reading, concept clarity, solved examples and practice questions. A formula sheet should be used for quick revision after understanding the chapter.