Study of Gas Laws Class 9 ICSE — Chemistry Chapter 7
Study of Gas Laws Class 9 ICSE is the focus of this page on Chemistry Chapter 7. The chapter explains how gases behave under changing pressure, volume and temperature. It defines what a gas is, then builds Boyle’s law, Charles’s law and the combined gas equation step by step. Each law comes with its formula, graphs, and a kinetic theory explanation. You also learn the Kelvin scale, S.T.P. conditions, and how moisture affects gas pressure.
ICSE Class 9 Chemistry Chapter 7 Study of Gas Laws PDF
What is inside Chapter 7
14 pages · 11 sections · 14 worked examples
| Section | Title |
|---|---|
| 7.1 | INTRODUCTION |
| 7.2 | BEHAVIOUR AND CHARACTERISTIC PROPERTIES OF GASES |
| 7.3 | MOLECULAR MOTION: RELATIONSHIP OF TEMPERATURE, PRESSURE AND VOLUME |
| 7.4 | THE GAS LAWS |
| 7.5 | PRESSURE AND VOLUME RELATIONSHIP IN GASES |
| 7.6 | TEMPERATURE-VOLUME RELATIONSHIP IN GASES |
| 7.8 | ABSOLUTE OR KELVIN SCALE OF TEMPERATURE |
| 7.9 | CONVERSION OF TEMPERATURE FROM CELSIUS SCALE TO KELVIN SCALE AND VICE-VERSA |
| 7.10 | THE GAS EQUATION (PRESSURE, VOLUME AND TEMPERATURE RELATIONSHIP) |
| 7.11 | STANDARD TEMPERATURE AND PRESSURE (S.T.P.) |
| 7.12 | THE EFFECT OF MOISTURE ON PRESSURE |
7.1 INTRODUCTION
Gas is the state of matter in which inter-particle attraction is weak and inter-particle space is large. The particles become completely free to move in the available space. A gas fills the entire vessel that contains it and takes the shape of that vessel.
All gases show uniform behaviour under similar conditions of temperature and pressure. This holds true irrespective of their chemical nature, colour, or odour. Because every gas follows the same physical rules, chemists describe their behaviour with a single set of laws.
This chapter follows the official CISCE syllabus for Class 9 Chemistry. The chapter builds these laws in a careful order. You start with the properties of gases, then learn the three variables that control their behaviour. After that come Boyle’s law, Charles’s law, and the gas equation that combines them.
7.2 BEHAVIOUR AND CHARACTERISTIC PROPERTIES OF GASES
Gases are made of tiny molecules in constant random motion. The kinetic molecular theory explains how these molecules produce the observed properties of gases. The theory gives eight characteristic properties.
- Composition: gases consist of tiny particles moving in all directions at all speeds. The particles are negligibly small compared with the volume the gas occupies.
- No fixed volume or shape: the force of attraction between particles is negligible. The particles move freely until the container walls stop them, so a gas takes the container’s shape.
- Pressure in all directions: moving particles collide with each other and with the container walls. These collisions exert pressure. At a given temperature the same number of molecules strikes a given area, so pressure is equal in all directions.
- High compressibility: large inter-molecular spaces let molecules move closer when pressure is applied. The volume shrinks.
- High expansibility: lowering pressure or raising temperature makes molecules move apart. The volume grows.
- Low density: few molecules per unit volume mean a gas has very low density.
- Diffusion: large inter-molecular spaces let gases mix. Diffusion is the gradual mixing of two substances in contact, caused by molecular motion.
- Liquefaction: cooling reduces molecular kinetic energy. Pressure then brings molecules closer, so the gas turns to liquid.
Diffusion is easy to observe. If you open a jar of chlorine in a large room, the smell spreads through the whole room within seconds. Chlorine is heavier than air, yet it does not stay near the floor. Its molecules collide with air molecules and scatter in every direction until the concentration is equal everywhere.
7.3 MOLECULAR MOTION: RELATIONSHIP OF TEMPERATURE, PRESSURE AND VOLUME
Gas molecules are always in motion, so they possess kinetic energy. The average kinetic energy of a molecule is directly proportional to its absolute temperature. As temperature rises, molecular motion speeds up. As temperature falls, it slows down.
Pressure comes from collisions. Molecules strike the walls of the container and exert a force on each unit area. That force per unit area is the pressure of the gas.
When temperature rises at constant volume, the molecules gain kinetic energy. They strike the walls with greater momentum and collide more often. Both effects raise the pressure. If a wall is movable, the higher pressure pushes it back and the gas expands.
Absolute zero is the temperature at which molecular motion completely ceases. On the Kelvin scale this is 0 K, equal to -273^{\circ}\text{C}. Temperatures below this are not possible. Temperature, pressure and volume are therefore linked through molecular motion. A change in one forces a change in the others. This link is the foundation for all the gas laws that follow.
7.4 THE GAS LAWS
Gas laws describe the behaviour of a gas under known conditions of pressure, volume and temperature. Three standard variables are used: volume V, pressure P, and temperature T. For a given mass of gas, a change in one variable changes the others.
Volume units matter for calculations. The S.I. unit is the cubic metre (\text{m}^3). Other common units are cubic centimetre (\text{cm}^3), millilitre (mL), cubic decimetre (\text{dm}^3) and litre (L).
| Unit | Equivalent |
|---|---|
| 1 m³ | 1000 dm³ = 1000 L |
| 1 L | 1000 mL = 1000 cm³ |
Pressure is the force a gas exerts per unit area on the walls of its container. The S.I. unit is the pascal (Pa), defined as a force of 1 newton on an area of 1 m². Other units include atmosphere (atm), cm Hg, mm Hg and torr.
| Unit | Equivalent |
|---|---|
| 1 atm | 76.0 cm Hg = 760 mm Hg = 760 torr |
Temperature is the degree of hotness of a gas. The Celsius scale allows negative values. Gas behaviour shows a lower limit at -273^{\circ}\text{C}, which leads to the Kelvin scale. Convert between them with \text{K} = ^{\circ}\text{C} + 273. This chapter sits within the wider Class 9 Chemistry book list, where the other chapters build on these gas laws.
7.5 PRESSURE AND VOLUME RELATIONSHIP IN GASES
Boyle studied the pressure-volume relationship with a syringe and piston. When weights press the piston down, the air volume shrinks. Removing the weights lets the piston rise and the volume grow again. The temperature stays constant throughout.
Boyle’s law: the volume of a given mass of a dry gas is inversely proportional to its pressure at constant temperature. If a gas occupies volume V_1 at pressure P_1, then the product stays constant: P_1V_1 = P_2V_2.
Three graphs verify the law. All are called isotherms, meaning at constant temperature.
- V versus 1/P: a straight line through the origin.
- V versus P: a curve.
- PV versus P: a straight line parallel to the pressure axis.
Kinetic theory explains why. Halving the volume doubles the number of molecules striking each unit area per second, so pressure doubles. The kinetic energy of the molecules stays the same because temperature is constant.
Significance: at constant temperature, density is directly proportional to pressure. At high altitudes, air pressure is low, so the air is less dense and carries less oxygen. That is why mountaineers carry oxygen cylinders.
One edge case: inflating a balloon increases both volume and pressure. This seems to break Boyle’s law. It does not, because the law applies to a fixed mass. Blowing more air into the balloon adds mass.
Worked pattern: a gas occupies 250 cm³ at 760 mm Hg. Find the pressure when it is compressed to 100 cm³ at constant temperature.
Step 1: Write the known values: P_1 = 760\ \text{mm Hg}, V_1 = 250\ \text{cm}^3, V_2 = 100\ \text{cm}^3.
Step 2: Apply P_1V_1 = P_2V_2.
P_2 = \frac{P_1V_1}{V_2} = \frac{760 \times 250}{100} = 1900\ \text{mm Hg}
Final answer: 1900 mm Hg, or 190 cm Hg.
7.6 TEMPERATURE-VOLUME RELATIONSHIP IN GASES
An inflated balloon left in bright sun bursts because the air inside expands as it heats. In general, raising temperature at constant pressure increases volume. Lowering temperature decreases it. Jacques Charles studied this relationship in 1787.
Charles’s law: at constant pressure, the volume of a given mass of a dry gas increases or decreases by 1/273 of its volume at 0^{\circ}\text{C} for each degree Celsius change. A simpler form uses absolute temperature: volume is directly proportional to Kelvin temperature at constant pressure.
The equation is \frac{V_1}{T_1} = \frac{V_2}{T_2}, where temperatures are in Kelvin. The graph of V versus T is a straight line. It is called an isobar, meaning constant pressure.
Kinetic theory explains the law. Heating makes molecules move faster. They strike the walls more often and more vigorously. To keep pressure constant, the volume must grow proportionally.
Significance: density decreases as temperature rises. That is why hot air balloons rise and are used for meteorological purposes.
The law also leads to absolute zero. Putting t = -273 into V = V_0(273 + t)/273 gives zero volume. In practice this is impossible. All gases liquefy or solidify first. It defines the lowest possible temperature.
Worked pattern: a gas occupies 300 cm³ at 27^{\circ}\text{C}. Find the temperature, in Celsius, at which it occupies 400 cm³ at constant pressure.
Step 1: Convert to Kelvin: T_1 = 27 + 273 = 300\ \text{K}.
Step 2: Apply \frac{V_1}{T_1} = \frac{V_2}{T_2}.
T_2 = \frac{V_2T_1}{V_1} = \frac{400 \times 300}{300} = 400\ \text{K}
Step 3: Convert back: 400 - 273 = 127^{\circ}\text{C}.
Final answer: 127 °C.
7.8 ABSOLUTE OR KELVIN SCALE OF TEMPERATURE
The Kelvin or absolute scale has its zero at -273.15^{\circ}\text{C}. For calculations, -273^{\circ}\text{C} is used. Each degree on the Kelvin scale is equal in size to one degree on the Celsius scale.
The scale is called absolute because its zero is the lowest possible temperature. No substance can be colder. The scale was built on the observation that gas volume falls to zero at -273^{\circ}\text{C}. Lord Kelvin proposed that this is the lowest temperature physically possible.
On this scale, water freezes at 273 K and boils at 373 K. These anchor values help you check conversions quickly. The Celsius and Kelvin scales differ only in where zero sits. The size of a degree is the same on both.
7.9 CONVERSION OF TEMPERATURE FROM CELSIUS SCALE TO KELVIN SCALE AND VICE-VERSA
The rule is simple. Add 273 to convert Celsius to Kelvin. Subtract 273 to convert Kelvin to Celsius.
- 0^{\circ}\text{C} = 273\ \text{K} (freezing point of water)
- 27^{\circ}\text{C} = 300\ \text{K}
- 373\ \text{K} = 100^{\circ}\text{C} (boiling point of water)
Worked pattern: convert 50^{\circ}\text{C} to Kelvin: 50 + 273 = 323\ \text{K}. Convert 250 K to Celsius: 250 - 273 = -23^{\circ}\text{C}.
Three notes matter. First, Kelvin values carry no degree sign. Second, one Kelvin degree equals one Celsius degree in size. Third, all Kelvin values are positive. That positivity is why gas laws use this scale.
Always convert temperatures to Kelvin before substituting into Charles’s law or the gas equation. Using Celsius values in a ratio gives a wrong answer, because the ratio changes when you shift the zero point.
7.10 THE GAS EQUATION (PRESSURE, VOLUME AND TEMPERATURE RELATIONSHIP)
Boyle’s law says volume varies inversely with pressure. Charles’s law says volume varies directly with absolute temperature. Combining them gives V \propto \frac{T}{P}, or \frac{PV}{T} = \text{constant}.
For a gas that changes from P_1, V_1, T_1 to P_2, V_2, T_2, the equation is \frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}. Temperature must be in Kelvin.
Worked pattern: a gas occupies 200 cm³ at 27^{\circ}\text{C} and 760 mm Hg. Find its volume at 127^{\circ}\text{C} and 1520 mm Hg.
Step 1: Convert temperatures: T_1 = 27 + 273 = 300\ \text{K}, T_2 = 127 + 273 = 400\ \text{K}.
Step 2: Apply the gas equation.
V_2 = \frac{P_1V_1T_2}{P_2T_1} = \frac{760 \times 200 \times 400}{1520 \times 300} = 133.3\ \text{cm}^3
Final answer: 133.3 cm³.
When one variable stays constant, the gas equation reduces to a simpler law. Constant temperature gives Boyle’s law. Constant pressure gives Charles’s law.
| Law | Constant condition | Formula | Graph shape |
|---|---|---|---|
| Boyle’s law | Temperature | P_1V_1 = P_2V_2 | V vs 1/P: straight line through origin |
| Charles’s law | Pressure | \frac{V_1}{T_1} = \frac{V_2}{T_2} | V vs T: straight line |
| Gas equation | None | \frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2} | Combines both laws |
7.11 STANDARD TEMPERATURE AND PRESSURE (S.T.P.)
Gas volume changes with temperature and pressure. To compare two gas volumes, they must be measured under the same conditions. The agreed reference conditions are called standard temperature and pressure (S.T.P.).
- Standard temperature: 0^{\circ}\text{C} = 273\ \text{K}
- Standard pressure: 760\ \text{mm Hg} = 76\ \text{cm Hg} = 1\ \text{atm}
Whenever a gas volume is stated, its temperature and pressure must also be given. Otherwise the number has no fixed meaning.
One useful note: increasing pressure shrinks a gas, while increasing temperature expands it. If both rise together, the two effects can cancel. The volume then stays constant.
Why does S.T.P. matter in calculations? When a problem says ‘at S.T.P.’, you plug in T = 273\ \text{K} and P = 760\ \text{mm Hg}. Many numericals in this chapter give one set of conditions at S.T.P. and ask you to convert a gas to or from it.
7.12 THE EFFECT OF MOISTURE ON PRESSURE
Some gases, like nitrogen and hydrogen, are collected over water. The collected gas is moist and contains water vapour. The water vapour exerts its own partial pressure, called aqueous tension.
The total pressure is the sum of the gas pressure and the water vapour pressure:
P_{\text{total}} = P_{\text{gas}} + P_{\text{water vapour}}
So the actual pressure of the dry gas is the total pressure minus the aqueous tension.
Worked pattern: suppose a moist gas is collected at 750 mm Hg. The aqueous tension at the collection temperature is 12 mm Hg. The dry gas pressure is 750 - 12 = 738\ \text{mm Hg}. This corrected value is what you substitute into the gas equation.
Students often skip this correction and use the measured total pressure as if the gas were dry. That error shifts every later calculation. Always subtract the aqueous tension before applying Boyle’s law or the gas equation.
What to practise from this chapter
The chapter closes with a descriptive exercise and three numerical sets. Work through them in this order.
- Descriptive exercise: 25 questions covering reasons, matching, fill in the blanks, and choose-the-correct-answer formats.
- Numericals based on Boyle’s law: 6 problems on pressure-volume changes at constant temperature.
- Numericals based on Charles’s law: 7 problems on temperature-volume changes, including temperature conversions.
- Final numerical set: 28 problems combining the gas equation, S.T.P. conversions, and aqueous tension corrections.
Pay attention to the graph questions in the descriptive set. You may be asked to identify which law a graph represents, or to plot P versus V, P versus 1/V, and PV versus P. The matching and fill-in-the-blank items test your recall of units, the Kelvin conversion, and the two laws’ formulas. For more practice material across subjects, browse the full Class 9 book collection.
Frequently asked questions
Why must temperature be in Kelvin when using Charles’s law and the gas equation?
Charles’s law and the gas equation use ratios of temperatures. Kelvin temperatures are all positive and start from absolute zero, so the ratios behave correctly. Celsius values can be negative, which breaks the proportionality. Always convert to Kelvin before substituting.
Why does inflating a balloon not violate Boyle’s law?
Boyle’s law applies only to a fixed mass of gas. Blowing into a balloon adds more air, so the mass increases. The law still holds for the gas already inside, but the changing mass means the simple P_1V_1 = P_2V_2 relation does not apply to the whole process.
What is aqueous tension and how do you correct gas pressure with it?
Aqueous tension is the partial pressure of water vapour in a moist gas collected over water. The dry gas pressure equals the total pressure minus the aqueous tension. You use this corrected pressure in the gas equation.
What are the S.T.P. values and why are gases compared at S.T.P.?
Standard temperature is 0^{\circ}\text{C} or 273 K. Standard pressure is 760 mm Hg, 76 cm Hg, or 1 atm. Gas volume changes with temperature and pressure, so a reference condition is needed to compare volumes fairly.
How do you convert a temperature from Celsius to Kelvin and back?
Add 273 to convert Celsius to Kelvin. Subtract 273 to convert Kelvin to Celsius. For example, 50^{\circ}\text{C} is 323 K, and 250 K is -23^{\circ}\text{C}.
Reference: ICSE Class 9 Chemistry textbook material, chapter 7: Study of Gas Laws — CISCE-aligned.