Chapter 13 States Of Matter

The Nature Of Gases kinetic = motion The energy an object has because of its motion is called kinetic energy. kinetic theory - All matter consists of tiny particles that are in constant motion. The particles in a gas are usually molecules or atoms. The particles in a gas are considered to be small, hard spheres with an insignificant volume. The motion of the particles in a gas is rapid, constant, and random. All collisions between particles in a gas are perfectly elastic.

The Nature Of Liquids The particles in a liquid are attracted to each other. Intermolecular Attractions -keep the particles in a liquid close together so liquids have a definite volume -reduce the amount of space between the particles in a liquid so liquids are much more dense than gases Increasing the pressure on a liquid has hardly any effect on its volume.

The Nature Of Solids Properties of solids reflect the orderly arrangement and fixed locations of their particles. In most solids, the atoms, ions, or molecules are packed tightly together. These solids are dense and not easy to compress. When you heat a solid, its particles vibrate more rapidly as their kinetic energy increases. The organization of particles within the solid breaks down, and eventually the solid melts. The melting point is the temperature at which a solid changes into a liquid. The melting and freezing points of a substance are at the same temperature. At that temperature, the liquid and solid phases are in equilibrium. Most solid substances are crystalline. In a crystal the particles are arranged in an orderly, repeating, three-dimensional pattern called a crystal lattice.

Non-Crystalline Solids The type of bonding that exists between particles in crystals determines their melting points. In general, ionic solids have high melting points because relatively strong forces hold them together. By contrast, molecular solids have relatively low melting points. Crystal Systems The shape of a crystal depends on the arrangement of the particles within it. The smallest group of particles within a crystal that retains the geometric shape of the crystal is known as a unit cell. A crystal lattice is a repeating array of any one of fourteen kinds of unit cells. Allotropes are two or more different molecular forms of the same element in the same physical state. Although allotropes are composed of atoms of the same element, they have different properties because their structures are different. Non-Crystalline Solids Not all solids are crystalline in form; some solids are amorphous. An amorphous solid lacks an ordered internal structure. Their atoms are randomly arranged. An example of amorphous solids are glasses. A glass is a transparent fusion product of inorganic substances that have cooled to a rigid state without crystallizing. Glasses are sometimes called supercooled liquids. The irregular internal structures of glasses are intermediate between those of a crystalline solid and those of a free-flowing liquid. Glasses do not melt at a definite temperature. Instead, they gradually soften when heated.

Gas Pressure 1 atm = 760 mm Hg = 101.3 kPa Gas pressure results from the force exerted by a gas per unit surface area of an object. It is the result of simultaneous collisions of billions of rapidly moving particles in a gas with an object. If there are no particles, there cannot be collisions. Consequently, there is no pressure. An empty space with no particles and no pressure is called a vacuum. Atmospheric pressure results from the collisions of atoms and molecules in air with objects. A barometer is a device that is used to measure atmospheric pressure. The SI unit of pressure is the pascal (Pa). It represents a very small amount of pressure. One standard atmosphere (atm) is the pressure required to support 760 mm of mercury in a mercury barometer at 25°C. The numerical relationship among the three units is given below. 1 atm = 760 mm Hg = 101.3 kPa

The Behavior Of Gases Gases are easily compressed, or squeezed into a smaller volume. Compressibility is a measure of how much the volume of matter decreases under pressure. The distance between particles in a gas is much greater than the distance between particles in a liquid or solid. Under pressure, the particles in a gas are forced closer together, or compressed. There are only a few nitrogen and oxygen molecules in this model of air. At room temperature, the distance between molecules in a container of air at standard pressure is about 10 times the diameter of a molecule.

The amount of gas, volume, and temperature are factors that affect gas pressure. increase gas increase in number of particles  increase in number of collisions  increase in gas pressure (gas reduced  pressure reduced) pressure (P) in kilopascals volume (V) in liters temperature (T) in kelvins As a gas is heated, the temperature increases and the average kinetic energy of the particles in the gas increases. Faster-moving particles impact the walls of their container with more energy.

Gas Laws P1 × V1 × V2= P2 V1/T1 = V2/T2 P1/T1 = P2/T2 Boyle’s Law: Pressure and Volume If the temperature is constant, as the pressure of a gas increases, the volume decreases. In turn, as the pressure decreases, the volume increases. Boyle’s law states that for a given mass of gas at constant temperature, the volume of the gas varies inversely with pressure. P1 × V1 × V2= P2 Charles’s Law: Temperature and Volume The amount of gas and the pressure are constant. As the temperature of an enclosed gas increases, the volume increases, if the pressure is constant. Charles’s law states that the volume of a fixed mass of gas is directly proportional to its Kelvin temperature if the pressure is kept constant. V1/T1 = V2/T2 Gay-Lussac’s Law: Pressure and Temperature As the temperature of an enclosed gas increases, the pressure increases, if the volume is constant. Gay-Lussac’s law states that the pressure of a gas is directly proportional to the Kelvin temperature if the volume remains constant. P1/T1 = P2/T2

The Combined Gas Law There is a single expression that combines Boyle’s, Charles’s, and Gay-Lussac’s laws. The combined gas law describes the relationship among the pressure, temperature, and volume of an enclosed gas. The combined gas law allows you to do calculations for situations in which only the amount of gas is constant. To calculate the number of moles of a contained gas requires an expression that contains the variable n. The number of moles of gas is directly proportional to the number of particles. The ideal gas constant (R) has the value 8.31 (L·kPa)/(K·mol). The gas law that includes all four variables—P, V, T, and n—is called the ideal gas law.

Ideal Gases and Real Gases An ideal gas is one that follows the gas laws at all conditions of P and T. At many conditions of T and P real gases behave like an ideal gas. The particles in a real gas do have volume, and there are attractions between the particles. Because of these attractions, a gas can condense, or even solidify, when it is compressed or cooled. The behavior of other real gases is similar, although lower temperatures and greater pressures may be required. Real gases differ most from an ideal gas at low temperatures and high pressures.

Daltons Law /Grahams Law Gas pressure results from collisions of particles in a gas with an object. If the number of particles increases in a given volume, more collisions occur. If the average kinetic energy of the particles increases, more collisions occur. In a mixture of gases the total pressure is the sum of the partial pressures of the gases. Dalton’s law of partial pressures states that, at constant volume and temperature, the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the component gases. Ptotal = P1 + P2 + P3 + … Diffusion is the tendency of molecules to move toward areas of lower concentration until the concentration is uniform throughout. During effusion, a gas escapes through a tiny hole in its container. Gases of lower molar mass diffuse and effuse faster than gases of higher molar mass. Graham’s law of effusion states that the rate of effusion of a gas is inversely proportional to the square root of the gas’s molar mass. This law can also be applied to the diffusion of gases.