About this EAMCET Engineering Chemistry practice test

40 of the 160 EAMCET marks sit in Chemistry, split fairly evenly across Physical, Organic and Inorganic — which means a strong score depends on being equally sharp at short calculations (moles, molarity, equilibrium) and equally solid on flat-out memorised facts (oxidation states, named reactions, IUPAC names). This practice set mirrors that split deliberately, so you're not over-preparing for one half while neglecting the other. Every answer includes the reasoning, whether that's a worked calculation or the rule being tested, so you always know why an answer is correct, not just what it is. Cycle in previous-year EAMCET questions once this feels comfortable to keep your recall sharp closer to the exam.

EAMCET Engineering Chemistry Practice Test sample questions

These starter questions help you launch a chemistry mock test quickly. Swap them with your own worksheet, notebook, or textbook questions any time.

1. 1. What is the maximum number of electrons that can occupy the M-shell (n = 3) of an atom?

2. 2. If the radius of the first Bohr orbit of a hydrogen atom is 0.529 Å, what is the radius of its third orbit (n = 3)?

3. 3. Among Li, Na, K and Rb, which element has the highest first ionization enthalpy?

4. 4. Among F, Cl, Br and I, which element has the highest electronegativity on the Pauling scale?

5. 5. What is the hybridization of the carbon atoms in ethyne (HC≡CH)?

6. 6. What is the molecular shape of SF6 according to VSEPR theory?

7. 7. A gas occupies 2.0 L at 1 atm and 300 K. What volume will it occupy at 2 atm at the same temperature?

8. 8. Calculate the pressure exerted by 2 moles of an ideal gas occupying 10 L at 300 K. (R = 0.0821 L atm K⁻¹ mol⁻¹)

9. 9. What is the value of ΔU for the isothermal expansion of an ideal gas?

10. 10. For a reaction with ΔH = -92 kJ/mol and Δn(gas) = -2 at 300 K, calculate ΔU. (R = 8.314 J K⁻¹ mol⁻¹)

11. 11. If the equilibrium constant Kc for a reaction is very large, what does this indicate about the position of equilibrium?

12. 12. What is the pH of a 0.01 M HCl solution, assuming complete dissociation?

13. 13. What is the oxidation state of chromium in potassium dichromate, K2Cr2O7?

14. 14. Calculate the standard cell potential of a Daniell cell given E°(Cu2+/Cu) = +0.34 V and E°(Zn2+/Zn) = -0.76 V.

15. 15. Calculate the molarity of a solution prepared by dissolving 4 g of NaOH (molar mass 40 g/mol) in enough water to make 500 mL of solution.

16. 16. How many atoms are present per unit cell in a face-centred cubic (fcc) lattice?

17. 17. A first-order reaction has a rate constant of 2 × 10⁻³ s⁻¹. Calculate its half-life.

18. 18. The rate of a reaction quadruples when the concentration of a single reactant is doubled. What is the order of the reaction with respect to that reactant?

19. 19. What is the most common oxidation state exhibited by the lanthanide elements?

20. 20. Which noble gas did Neil Bartlett first convert into a chemical compound, XePtF6?

21. 21. What is the coordination number of the central metal ion in [Co(NH3)6]3+?

22. 22. What is the oxidation state of nickel in the complex [Ni(CO)4]?

23. 23. What is the IUPAC name of the hydrocarbon CH3-CH2-CH2-CH3?

24. 24. Which catalyst is used to convert an alkyne into a cis-alkene by controlled hydrogenation?

25. 25. Which chemical test is used to distinguish primary, secondary, and tertiary alcohols based on differing reaction rates?

26. 26. Which test gives a violet or purple coloration characteristic of phenols but not alcohols?

27. 27. Which test is used to detect aldehydes by the formation of a silver mirror on the test tube?

28. 28. Arrange formic acid, acetic acid, and benzoic acid in decreasing order of acidity.

29. 29. What type of chemical bond links amino acids together in a protein chain?

30. 30. Which class of drugs, including aspirin and paracetamol, relieves pain without affecting consciousness and without causing addiction?

Answer key & quick explanations

Short answers for the sample questions above. Use this to self-check before generating a fresh AI-built mock test.

1. 1. What is the maximum number of electrons that can occupy the M-shell (n = 3) of an atom?

   18

   The capacity of any shell is given by 2n². For n = 3, this works out to 2 × 9 = 18 electrons, spread across the 3s, 3p and 3d subshells.

2. 2. If the radius of the first Bohr orbit of a hydrogen atom is 0.529 Å, what is the radius of its third orbit (n = 3)?

   4.761 Å

   For hydrogen, orbit radius scales as rn = r1 × n² since Z = 1. Substituting n = 3 gives 0.529 × 9 = 4.761 Å.

3. 3. Among Li, Na, K and Rb, which element has the highest first ionization enthalpy?

   Lithium (Li)

   Ionization enthalpy falls as you move down a group because atomic radius increases and the outermost electron is held less tightly. Lithium, being the smallest of the four, needs the most energy to lose an electron.

4. 4. Among F, Cl, Br and I, which element has the highest electronegativity on the Pauling scale?

   Fluorine (F)

   Electronegativity decreases down a group as atomic size and shielding increase. Fluorine's very small atomic radius lets its nucleus pull bonding electrons most strongly, giving it a Pauling value of about 3.98, the highest of any element.

5. 5. What is the hybridization of the carbon atoms in ethyne (HC≡CH)?

   sp

   Each carbon in ethyne forms one sigma bond to hydrogen and one sigma bond to the other carbon, plus two pi bonds using unhybridized p orbitals. This leaves only two hybrid orbitals needed per carbon, so sp hybridization results, giving the linear 180° geometry.

6. 6. What is the molecular shape of SF6 according to VSEPR theory?

   Octahedral

   Sulfur forms six bonding pairs with fluorine and has no lone pairs left over. Six regions of electron density around a central atom arrange themselves at 90° to one another, producing an octahedral shape.

7. 7. A gas occupies 2.0 L at 1 atm and 300 K. What volume will it occupy at 2 atm at the same temperature?

   1.0 L

   Temperature is unchanged, so Boyle's law applies: P1V1 = P2V2. Solving gives V2 = (1 × 2.0)/2 = 1.0 L, showing volume is inversely proportional to pressure.

8. 8. Calculate the pressure exerted by 2 moles of an ideal gas occupying 10 L at 300 K. (R = 0.0821 L atm K⁻¹ mol⁻¹)

   ≈ 4.93 atm

   Rearranging the ideal gas equation PV = nRT gives P = nRT/V. Plugging in the values, P = (2 × 0.0821 × 300)/10 = 49.26/10 ≈ 4.93 atm.

9. 9. What is the value of ΔU for the isothermal expansion of an ideal gas?

   Zero

   The internal energy of an ideal gas is a function of temperature alone. Because the process is isothermal, temperature does not change, so ΔU = 0 no matter how much work is done or how much heat flows.

10. 10. For a reaction with ΔH = -92 kJ/mol and Δn(gas) = -2 at 300 K, calculate ΔU. (R = 8.314 J K⁻¹ mol⁻¹)

    ≈ -87.0 kJ/mol

    Using ΔH = ΔU + Δn(gas)RT, rearrange to ΔU = ΔH − Δn(gas)RT. Here Δn(gas)RT = -2 × 8.314 × 300 = -4988.4 J ≈ -4.99 kJ, so ΔU = -92 - (-4.99) ≈ -87.0 kJ/mol.

11. 11. If the equilibrium constant Kc for a reaction is very large, what does this indicate about the position of equilibrium?

    The equilibrium lies mostly toward the products

    Kc is the ratio of product to reactant concentrations at equilibrium. A large numerical value only arises when product concentrations far outweigh reactant concentrations, meaning the forward reaction has gone nearly to completion.

12. 12. What is the pH of a 0.01 M HCl solution, assuming complete dissociation?

    pH = 2

    HCl is a strong acid and ionizes fully in water, so [H+] equals the stated concentration, 0.01 M or 10⁻² M. Taking the negative log gives pH = -log(10⁻²) = 2.

13. 13. What is the oxidation state of chromium in potassium dichromate, K2Cr2O7?

    +6

    Assign K as +1 and O as -2 and let chromium be x. The compound is neutral, so 2(+1) + 2x + 7(-2) = 0, which gives 2x = 12 and x = +6.

14. 14. Calculate the standard cell potential of a Daniell cell given E°(Cu2+/Cu) = +0.34 V and E°(Zn2+/Zn) = -0.76 V.

    +1.10 V

    In a Daniell cell, copper is reduced at the cathode and zinc is oxidized at the anode. Cell potential is found from E°cell = E°cathode − E°anode = 0.34 − (-0.76) = 1.10 V.

15. 15. Calculate the molarity of a solution prepared by dissolving 4 g of NaOH (molar mass 40 g/mol) in enough water to make 500 mL of solution.

    0.2 M

    First find moles of NaOH: 4 g ÷ 40 g/mol = 0.1 mol. Molarity is moles divided by volume in litres, so 0.1 mol ÷ 0.5 L = 0.2 M.

16. 16. How many atoms are present per unit cell in a face-centred cubic (fcc) lattice?

    4

    An fcc cell has 8 corner atoms, each shared among 8 cells (contributing 8 × 1/8 = 1 atom), plus 6 face-centred atoms, each shared between 2 cells (contributing 6 × 1/2 = 3 atoms). Adding these gives 1 + 3 = 4 atoms per cell.

17. 17. A first-order reaction has a rate constant of 2 × 10⁻³ s⁻¹. Calculate its half-life.

    ≈ 346.5 s

    For any first-order process, half-life is independent of starting concentration and follows t½ = 0.693/k. Substituting k = 2 × 10⁻³ s⁻¹ gives t½ = 0.693/0.002 ≈ 346.5 s.

18. 18. The rate of a reaction quadruples when the concentration of a single reactant is doubled. What is the order of the reaction with respect to that reactant?

    Second order

    Rate depends on concentration as rate = k[A]ⁿ. Doubling [A] and seeing the rate rise fourfold means 2ⁿ = 4, which is only satisfied when n = 2, so the reaction is second order in that reactant.

19. 19. What is the most common oxidation state exhibited by the lanthanide elements?

    +3

    Lanthanides typically lose their two 6s electrons along with one inner d or f electron, settling into a stable +3 state. Although a few members also show +2 or +4 states, +3 dominates across the entire series.

20. 20. Which noble gas did Neil Bartlett first convert into a chemical compound, XePtF6?

    Xenon

    In 1962, Bartlett reacted xenon gas with platinum hexafluoride to obtain xenon hexafluoroplatinate, XePtF6. This landmark result overturned the belief that noble gases could never form compounds.

21. 21. What is the coordination number of the central metal ion in [Co(NH3)6]3+?

    6

    Six ammonia molecules each donate a lone pair of electrons to the cobalt(III) centre, so six donor atoms surround the metal. This gives a coordination number of 6 and an octahedral arrangement.

22. 22. What is the oxidation state of nickel in the complex [Ni(CO)4]?

    0

    Carbon monoxide is a neutral ligand and the overall complex carries no charge. Since none of the ligands contribute charge, the nickel atom itself must be in the zero oxidation state.

23. 23. What is the IUPAC name of the hydrocarbon CH3-CH2-CH2-CH3?

    Butane

    This molecule is a straight, saturated chain of exactly four carbon atoms with no branching or substituents. Following IUPAC alkane nomenclature, a four-carbon chain is named butane.

24. 24. Which catalyst is used to convert an alkyne into a cis-alkene by controlled hydrogenation?

    Lindlar's catalyst (Pd on CaCO3, poisoned with lead and quinoline)

    Lindlar's catalyst is a partially deactivated palladium surface that adds both hydrogen atoms to the same face of the triple bond. This syn addition halts the reaction at the alkene stage and yields the cis (Z) isomer specifically.

25. 25. Which chemical test is used to distinguish primary, secondary, and tertiary alcohols based on differing reaction rates?

    Lucas test

    Lucas reagent is a mixture of concentrated HCl and anhydrous ZnCl2. Tertiary alcohols react instantly and turn cloudy, secondary alcohols take a few minutes to show turbidity, while primary alcohols show no visible change at room temperature.

26. 26. Which test gives a violet or purple coloration characteristic of phenols but not alcohols?

    Neutral ferric chloride (FeCl3) test

    Phenols react with neutral FeCl3 solution to form a colored iron-phenolate complex, usually appearing violet, blue or purple. Ordinary alcohols lack the aromatic -OH needed for this complex, so they give no such color change.

27. 27. Which test is used to detect aldehydes by the formation of a silver mirror on the test tube?

    Tollens' test

    Tollens' reagent contains a diammine silver(I) complex that oxidizes the aldehyde group to a carboxylate while silver ions are reduced to metallic silver. The freshly deposited silver coats the glass, producing the familiar mirror effect that ketones do not give.

28. 28. Arrange formic acid, acetic acid, and benzoic acid in decreasing order of acidity.

    Formic acid > Benzoic acid > Acetic acid

    Formic acid has only a hydrogen atom attached to the carboxyl carbon, so there is no electron-donating group to weaken its acidity, making it the strongest of the three. Benzoic acid's ring is more electron-withdrawing than the electron-donating methyl group in acetic acid, placing it in between.

29. 29. What type of chemical bond links amino acids together in a protein chain?

    Peptide bond (amide linkage)

    A peptide bond forms when the carboxyl group of one amino acid reacts with the amino group of the next, releasing a molecule of water in a condensation reaction. Repeating this linkage along many amino acid units builds up the polypeptide backbone of a protein.

30. 30. Which class of drugs, including aspirin and paracetamol, relieves pain without affecting consciousness and without causing addiction?

    Non-narcotic analgesics

    Non-narcotic analgesics work mainly by inhibiting prostaglandin synthesis at the site of pain rather than acting on the central nervous system. This is why aspirin and paracetamol control pain without impairing consciousness or creating the dependence seen with narcotic analgesics such as morphine.

Curriculum Mapping & Learning Guide

Use this breakdown to identify which skills each question tests and guide post-test review.

EAMCET Engineering Chemistry units covered

1. Chapter 1: Atomic Structure
2. Chapter 2: Classification of Elements and Periodicity
3. Chapter 3: Chemical Bonding and Molecular Structure
4. Chapter 4: States of Matter (Gases, Liquids, Solids)
5. Chapter 5: Thermodynamics
6. Chapter 6: Equilibrium
7. Chapter 7: Redox Reactions and Electrochemistry
8. Chapter 8: Solutions and Solid State
9. Chapter 9: Chemical Kinetics
10. Chapter 10: p-Block, d-Block and f-Block Elements
11. Chapter 11: Coordination Compounds
12. Chapter 12: Hydrocarbons
13. Chapter 13: Alcohols, Phenols and Ethers
14. Chapter 14: Aldehydes, Ketones and Carboxylic Acids
15. Chapter 15: Biomolecules and Polymers
16. Chapter 16: Chemistry in Everyday Life

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