Is States of Matter: Gases and Liquids part of the EAPCET Chemistry syllabus?

Yes. States of Matter: Gases and Liquids is taught in Intermediate first year Chemistry, and the EAPCET Engineering syllabus is based on the Intermediate first- and second-year syllabus. Check the official notification each year for any changes.

How many Chemistry questions are there in EAPCET?

40 of the 160 questions in the Engineering paper are Chemistry, spread across 24 chapters.

What mistakes do students make in States of Matter: Gases and Liquids?

Using temperature in °C in PV = nRT or Charles’ law — always convert to kelvin by adding 273. Mixing R values — use 0.0821 L atm K⁻¹ mol⁻¹ with atm and litres, and 8.314 J K⁻¹ mol⁻¹ with SI units. Forgetting to square the rate ratio in Graham’s law — rates vary with the square root of M, so a rate ratio of 1/2 means a mass ratio of 4. Thinking all real gases have Z 1 because the b (size) term dominates.

EAPCET Chemistry · Intermediate first year

States of Matter: Gases and Liquids

States of Matter covers the gas laws, the ideal gas equation, Dalton’s and Graham’s laws, the kinetic theory of gases and real-gas behaviour.

Subject: Chemistry
Taught in: Intermediate first year
Chemistry in the paper: 40 of 160 questions

Key formulas and results

$PV = nRT$
Dalton: total pressure = sum of partial pressures
Graham: rate of diffusion $\propto \frac{1}{\sqrt{M}}$
van der Waals: $\left(P + \frac{an^2}{V^2}\right)(V - nb) = nRT$
Compressibility factor $Z = \frac{PV}{nRT}$ ( $Z = 1$ for an ideal gas)
Values of R: $0.0821\ \text{L atm K}^{-1}\,\text{mol}^{-1} = 8.314\ \text{J K}^{-1}\,\text{mol}^{-1} = 1.987\ (\approx 2)\ \text{cal K}^{-1}\,\text{mol}^{-1} = 62.36\ \text{L torr K}^{-1}\,\text{mol}^{-1}$
Partial pressure of a gas in a mixture: $p_i = x_i \times P_{\text{total}}$
Average molar mass of a gas mixture: $M_{\text{avg}} = \frac{n_1M_1 + n_2M_2 + \dots}{n_1 + n_2 + \dots}$

What this chapter covers

Boyle’s, Charles’ and Avogadro’s laws
Ideal gas equation
Dalton’s law of partial pressures
Graham’s law of diffusion
Kinetic molecular theory
Real gases and the van der Waals equation

How it is usually asked

Gas law calculations
Comparing diffusion rates
Real-gas behaviour and Z
Partial pressures in mixtures

Understanding States of Matter: Gases and Liquids

Gases are described by four variables: pressure, volume, temperature and amount. Boyle’s law (PV constant at fixed T), Charles’ law (V ∝ T at fixed P) and Avogadro’s law (V ∝ n) combine into the ideal gas equation PV = nRT. Temperature must always be in kelvin. Because it also gives PM = dRT, the same equation lets you find molar mass or density of a gas from measured conditions.

In a mixture, each gas behaves as if it were alone: its partial pressure equals its mole fraction times the total pressure (Dalton’s law). Graham’s law says lighter gases diffuse and effuse faster, with rate inversely proportional to the square root of molar mass. Kinetic theory explains all this: average kinetic energy depends only on temperature, so at the same T lighter molecules simply move faster. RMS speed is √(3RT/M), larger than the average and most probable speeds.

Real gases deviate from ideal behaviour at high pressure and low temperature because molecules occupy volume and attract each other. The van der Waals equation corrects pressure with a (attraction) and volume with b (molecular size). The compressibility factor Z = PV/nRT is 1 for an ideal gas; Z < 1 means attractions dominate and Z > 1 means repulsion or size dominates. H₂ and He show Z > 1 at ordinary temperatures.

Worked example

EAPCET-style question

Under identical conditions, a gas X diffuses at half the rate of oxygen. The molar mass of X is:

(1)16 g mol⁻¹
(2)64 g mol⁻¹
(3)128 g mol⁻¹
(4)8 g mol⁻¹

Solution

Graham’s law: $\frac{r_X}{r_{\mathrm{O_2}}} = \sqrt{\frac{M_{\mathrm{O_2}}}{M_X}}$ .
Given $\frac{r_X}{r_{\mathrm{O_2}}} = \frac{1}{2}$ and $M_{\mathrm{O_2}} = 32\ \text{g mol}^{-1}$ .
Square both sides: $\frac{1}{4} = \frac{32}{M_X}$ .
$M_X = 32 \times 4 = 128\ \text{g mol}^{-1}$ .
64 comes from forgetting to square, and 8 from inverting the ratio — both are common slips.

Answer: 128 g mol⁻¹

Common mistakes to avoid

Using temperature in °C in PV = nRT or Charles’ law — always convert to kelvin by adding 273.
Mixing R values — use 0.0821 L atm K⁻¹ mol⁻¹ with atm and litres, and 8.314 J K⁻¹ mol⁻¹ with SI units.
Forgetting to square the rate ratio in Graham’s law — rates vary with the square root of M, so a rate ratio of 1/2 means a mass ratio of 4.
Thinking all real gases have Z < 1 — at ordinary temperatures H₂ and He have Z > 1 because the b (size) term dominates.

How to prepare it

Match the value of R to your units before calculating — it is the most common slip in this chapter.

Once the formulas feel familiar, test them the way the exam does: practise States of Matter: Gases and Liquids questions chapter-wise, then attempt a full previous year paper to see whether it holds up under a 3-hour timer.

Related Chemistry chapters

See every chapter in the EAPCET Chemistry syllabus.

Practise States of Matter: Gases and Liquids chapter-wise

Chapter-wise quizzes show exactly which topics in this chapter you still get wrong — before a full paper does.

Frequently asked questions

Is States of Matter: Gases and Liquids part of the EAPCET Chemistry syllabus?: Yes. States of Matter: Gases and Liquids is taught in Intermediate first year Chemistry, and the EAPCET Engineering syllabus is based on the Intermediate first- and second-year syllabus. Check the official notification each year for any changes.
How should I prepare States of Matter: Gases and Liquids for EAPCET?: Match the value of R to your units before calculating — it is the most common slip in this chapter.
How many Chemistry questions are there in EAPCET?: 40 of the 160 questions in the Engineering paper are Chemistry, spread across 24 chapters.
What mistakes do students make in States of Matter: Gases and Liquids?: Using temperature in °C in PV = nRT or Charles’ law — always convert to kelvin by adding 273. Mixing R values — use 0.0821 L atm K⁻¹ mol⁻¹ with atm and litres, and 8.314 J K⁻¹ mol⁻¹ with SI units. Forgetting to square the rate ratio in Graham’s law — rates vary with the square root of M, so a rate ratio of 1/2 means a mass ratio of 4. Thinking all real gases have Z < 1 — at ordinary temperatures H₂ and He have Z > 1 because the b (size) term dominates.