THE KINETIC THEORY OF MATTER

The kinetic theory of matter postulates that particles of a substance are continually moving and so possess kinetic energy.

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• STATES OF MATTER

Matter exists in three physical state – solid, liquid and gas. Solid has a definite volume and shape; liquid has a definite volume but no definite volume (takes the shape of its container); and a gas has neither a definite volume nor shape (fills the entire container).

• CHANGES IN STATE OF MATTER

Changes in the state of matter usually occur due to the application of heat / increase in temperature. These changes take place under the following conditions:

• MELTING – A change from solid to liquid due to the application of heat, usually at a certain temperature called the melting point.
• SUBLIMATION – A direct change from solid to gas, without going through the liquid phase.
• EVAPORATION – A change from liquid to gas. It occurs at all temperatures; rate of evaporation increases as temperature increases.
• FREEZING – A change from liquid to solid due to the cooling effect experienced by the liquid. This usually occurs at a certain temperature called the freezing point.
• CONDENSATION – A change in a substance from its gaseous state / vapour to liquid, usually caused by a cooling effect.
• BOILING – Occurs when a liquid is heated to a certain temperature called the boiling point (a temperature at which the saturated vapour pressure of the liquid equals the atmospheric pressure.

• PHENOMENA SUPPORTING THE KINEIC THEORY OF MATTER

• BROWNIAN MOTION – Observed by Brown (a botanist, in 1827). He observed that particles of a substance are always in constant state of random motion.
• DIFFUSION – The movement of solute particles through a medium, from a region of higher concentration to a region of lower concentration, until it is evenly distributed. It is fastest in gas, faster in liquid and fast in solid.
• OSMOSIS – The movement of water molecules through a semi permeable membrane, from a region of higher concentration to a region of lower concentration.

• KINETIC THEORY OF GASES

• The gas molecules move randomly in straight lines, colliding with one another and with the walls of the container.
• The collisions of gas molecules are perfectly elastic.
• The actual volume occupied by the gas molecules themselves is negligible.
• The cohesive forces between the gas molecules are negligible.
• The temperature of the gas is a measure of the average kinetic energy of the gas particles.

NB: The kinetic theory of gases is only applicable for an ideal or perfect gas.

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• GAS LAWS

• BOYLE’S LAW – It states that the volume of a given mass of a gas is inversely proportional to its pressure, provided that the temperature remains constant (i.e P₁V₁ = P₂V₂).
• CHARLES’ LAW – It states that the volume of a given mass of gas is directly proportional to its absolute temperature in kelvin, provided that pressure remains constant (i.e V₁ / T₁ = V₂ /T₂).
• GAY LUSSAC’S PRESSURE LAW – It states that the pressure of a given mass of gas is directly proportional to its absolute temperature in kelvin, provided that volume remains constant (i.e P₁ / T₁ = P₂ /T₂).
• GAY LUSSAC’S LAW OF COMBINING VOLUMES – It states that when gases react, they do so in volumes which are in simple ratio to one another, and to the volume of the product(s) formed if gaseous, provided the temperature and pressure remain constant.
• AVOGADRO’S LAW – It state that under the same conditions of temperature and pressure, equal volumes of all gases contain an equal number of molecules.
• DALTON’S LAW OF PARTIAL PRESSURE – It states that in a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of the individual gases ( P_T = P₁ + P₂  + P₃ ….. + P_n).
• GRAHAM’S LAW OF DIFFUSION – It states that at constant temperature and pressure, the rate of diffusion of a gas is inversely proportional to the square root of its density. (i.e R₁ / R₂ = ✓d₂ / ✓d₁). It could also be stated as; at constant temperature and pressure and pressure, the rate of diffusion of a gas is inversely proportional to the square root of it molecular mass (i.e R ₁ / R ₂ = ✓M₂ / ✓M₁).
• GENERAL GAS EQUATION – A combination of Charles’, Boyle’s and Pressure law. It shows the relationship between pressure, temperature and volume of a gas (i.e (P₁ × V₁) / T₁ = (P₂ × V₂) / T₂).
• IDEAL GAS EQUATION – All gases obey an equation of state known as the ideal gas law: PV = nRT.  n is the number of moles of the gas and R is the molar gas constant (R = 0.082057 atm dm³ mol^-1 K^-1 or 8.314 J mol^-1 K^-1. P is pressure, V is volume and T is temperature.

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• RELATIONSHIP BETWEEN VAPOUR DENSITY AND RELATIVE MOLECULAR MASS

Relative Molecular Mass (M) = 2 × Vapour Density (VD)

Vapour Density (VD) = ½ × Relative Molecular Mass (M)

• Vapour density is the mass of a certain volume of a gas or vapour compared to the mass of the same volume of hydrogen gas at a particular temperature and pressure.
• Relative molecular mass (M) refers to the number of times, one molecule of a substance is heavier than one-twelfth the mass of one atom of carbon-12, often represented as the sum of the atomic masses of the elements in the molecule.

Since hydrogen gas (H₂) is the reference and its relative molecular mass is 2, the relationship between vapour density and relative molecular mass becomes M = 2×VD.

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  GRAPHICAL ILLUSTRUATIONS OF THE GAS LAWS

• MOLAR VOLUME – The molar volume of any gas is the volume occupied by one mole of that gas at s.t.p. and is numerically equal to 22.4dm³.

NB – At standard temperature and pressure,

• T = 0^oC or 273K   |   P = 760mmHg or 1.01 × 10⁵Nm^-2