www.tiwariacademy.in UNIT 5 : STATES OF MATTER CONCEPT WISE HANDOUTS KEY CONCEPTS : 1. Intermolecular Forces 2. Gas Laws 3. Behaviour of gases Concept 1. INTERMOLECULAR FORCES Intermolecular forces- forces of attraction and repulsion between molecules that hold molecules, ions, and atoms together. Intramolecular forces - forces of chemical bonds within a molecule Intermolecular vs Intramolecular 41 kj to vaporize 1 mole of water (inter) 930 kj to break all O-H bonds in 1 mole of water (intra) Generally, INTERmolecular forces are much weaker than INTRAmolecular forces. Types of intermolecular forces Intermolecular forces are the forces of attraction and repulsion between interacting particles (atoms and molecules). This term does not include the electrostatic forces that exist between the two oppositely charged ions ionic bond and the forces that hold atoms of a molecule together i.e., covalent bond. Van der Waals forces: the attraction of intermolecular forces between molecules. There are two kinds of Van der Waals forces: weak London Dispersion Forces and stronger dipole-dipole forces. London dispersion forces, named after the German-American physicist Fritz London, are weak intermolecular forces that arise from the interactive forces between instantaneous multipoles in molecules without permanent multipole moments. These forces dominate the interaction of non-polar molecules, and also play a less significant role in van der Waals forces than molecules containing permanent dipoles or ionized molecules.
www.tiwariacademy.in London dispersion forces are also known as dispersion forces, London forces, or instantaneous dipole induced dipole forces. They increase with the molar mass, causing a higher boiling point especially for the halogen group. Dipole dipole forces Dipole-dipole Interactions are stronger intermolecular forces than Dispersion forces & occur between molecules that have permanent net dipoles (polar molecules). The partial positive charge on one molecule is electro statically attracted to the partial negative charge on a neighboring molecule. Dipole induced Dipole An induced dipole is caused when a molecule that would not by itself have a dipole moment is brought close to a molecule with a dipole moment. (That is, the dipole "pushes" the electron cloud of the symmetric molecule. A dipole-induced dipole attraction is a weak attraction that results when a polar molecule induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.
Hydrogen bond www.tiwariacademy.in The hydrogen bond is really a special case of dipole forces. A hydrogen bond is the attractive force between the hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different molecule. Usually the electronegative atom is oxygen, nitrogen, or fluorine, which has a partial negative charge. The hydrogen then has the partial positive charge. In molecules containing N-H, O-H or F-H bonds, the large difference in electronegativity between the H atom and the N, O or F atom leads to a highly polar covalent bond (i.e., a bond dipole. ELEMENT ELECTRONEGATIVITY VALUE H 2.1 N 3.0 O 3.5 F 4.1 INTERMOLECULAR HYDROGEN BOND INTRAMOLECULAR HYDROGEN BOND
www.tiwariacademy.in Concept 2. GAS LAWS BOYLE S LAW One of the most amazing things about gases is that, despite wide differences in chemical properties, all the gases more or less obey the gas laws. The gas laws deal with how gases behave with respect to pressure, volume, temperature, and amount. Boyle's Law (sometimes referred to as the Boyle Mariotte law) states that the absolute pressure and volume of a given mass of a confined gas are inversely proportional, if the temperature remains unchanged within a closed system. other halves). CHARLE S LAW The law itself can be stated as follows: for a fixed amount of an ideal gas kept at a fixed temperature, P (pressure) and V (volume) are inversely proportional (when one doubles, the Pressure α 1 / volume P = pressure in N/m 2 ; V=volume in dm 3 (litres) ; k=constant PV=k P 1 V 1 =P 2 V 2 French chemist Jacques Charles discovered that the volume of a gas at constant pressure changes with temperature. The Law states that at constant pressure, the volume of a fixed number of particles of gas is directly proportional to the absolute (Kelvin) temperature, mathematically expressed as: V = k T V = Volume k = Charles Law constant of Proportionality T = Temperature in Kelvins Raising the temperature of a gas causes the gas to fill a greater volume as long as pressure remains constant. Gases expand at a constant rate as temperature increases, and the rate of expansion is similar for all gases. In a sample with volume V 1 & temperature T 1, changing either volume or temperature converts these variables to V 2 and T 2.
V 1 / T 1 = k = V 2 / T 2 Therefore: V 1 T 2 = V 2 T 1 Charles saw a linear relationship between the volume and temperature of a gas. Extrapolating backwards, he found that the point where a gas would have no volume would be -273 degrees Celsius. Since that's as cold as he thought things could ever get, that originated the idea of absolute zero. www.tiwariacademy.in As it turns out, 0 K is the lowest temperature that can be achieved, but not because of Charles law. It actually has to do with the fact that at absolute zero, molecules have the smallest amount of energy possible, and have ceased moving back and forth entirely. It's sometimes thought that molecules don't move at all at absolute zero - but there's still some vibrational motion in bonds even at this temp. GAY-LUSSAC'S LAW Gay-Lussac's law is an ideal gas law where at constant volume, the pressure of an ideal gas is directly proportional to its absolute temperature. In Gay-Lussac s law the pressure exerted by a gas is directly related to the Kelvin temperature. V and n are constant.
www.tiwariacademy.in AVAGADROS LAW.Avogadro's Law proposed by Amedeo Avogadro states that under equal conditions of temperature and pressure, equal volumes of gases contain an equal number of molecules. This hypothesis was not acknowledged in Avogadro's lifetime and it wasn't until Stanislao Cannizzaro, in 1860, demonstrated that it was the solution to the problem of atomic and molecular weights that Avogadro's Law became widely accepted. 10 23 The volume of a gas is directly proportional its number of moles (n), regardless of the IDEAL GAS EQUATION identity of the gas V n The ideal gas law is the equation of state of a hypothetical ideal gas. It is a good approximation to the behaviour of many gases under many conditions, although it has several limitations. Gases can described in terms of four variables: pressure (P), volume (V), temperature (T), and the amount of gas (n). There are five relationships between pairs of these variables in which two of the variables were allowed to change while the other two were held constant. P α n (T and V constant) Boyle's law: P α 1/V (T and n constant) Gay Lusac s law: P α T (V and n constant) Charles' law: V α T (P and n constant) Avogadro's hypothesis: V α n (P and T constant) Each of these relationships is a special case of a more general relationship known as the ideal gas equation. In this equation, R is a proportionality constant known as the ideal gas constant and T is the absolute temperature. The value of R depends on the units used to express the four variables P, V, n, and T. By convention, most chemists use the following set of units. P: atmospheres T: kelvin V: liters n: moles
www.tiwariacademy.in DALTONS LAW OF PARTIAL PRESSURE Dalton's Law of Partial Pressures, or Dalton's Law, states that the total pressure of a gas in a container is the sum of the partial pressures of the individual gases in the container. Here is a worked example problem showing how to use Dalton's Law to calculate the pressure of a gas. Dalton's Law of Partial Pressures is a gas law that can be stated: P total = P 1 + P 2 + P 3 +... P n where P 1, P 2, P 3, P n are the partial pressures of the individual gases in the mixture. PARTIAL PRESSURE MOLE FRACTION The mole fraction of an individual gas component in an ideal gas mixture can be expressed in terms of the component's partial pressure or the moles of the component: and the partial pressure of an individual gas component in an ideal gas can be obtained using this expression: where: x i P i n i n P = mole fraction of any individual gas component in a gas mixture = partial pressure of any individual gas component in a gas mixture = moles of any individual gas component in a gas mixture = total moles of the gas mixture = pressure of the gas mixture The mole fraction of a gas component in a gas mixture is equal to the volumetric fraction of that component in a gas mixture MOLAR MASS AND GAS DENSITIES
KINETIC MOLECULAR THEORY www.tiwariacademy.in A gas consists of very small particles, each of which has a mass. An inflated basketball weighs more than a deflated basketball. The distances separating gas particles are relatively large. The volume of the gas particles is assumed to be zero because it is negligible compared with the total volume in which the gas is contained. Gas particles are in constant, rapid, random motion. Gases immediately fill a container and quickly diffuse from one area to another. Collisions of gas particles with each other or with the walls of the container are perfectly elastic. Unlike bouncing balls, no energy of motion is lost. The average kinetic energy of gas particles depends only on the temperature of the gas.the kinetic energy of gas molecules is proportional to their temperature in Kelvins -- a good description. KE = mv 2 / 2 High T Higher KE Low T Lower KE Gas particles exert no force on one another. Attractive forces between gas particles is assumed to be zero. Gas particles do not slow down and condense into a liquid because they exert only very weak attractive forces upon each other. Gas molecules don t interact with one another. Concept 3. BEHAVIOR OF GASES The behavior of real gases usually agrees with the predictions of the ideal gas equation to within 5% at normal temperatures and pressures. At low temperatures or high pressures, real gases deviate significantly from ideal gas behavior. In 1873, while searching for a way to link the behavior of liquids and gases, the Dutch physicist Johannes van der Waals developed an explanation for these deviations and an equation that was able to fit the behavior of real gases over a much wider range of pressures. Van der Waals realized that two of the assumptions of the kinetic molecular theory were questionable. The kinetic theory assumes that gas particles occupy a negligible fraction of the total volume of the gas. It also assumes that the force of attraction between gas molecules is zero.
Volume Correction www.tiwariacademy.in The actual volume free to move in is less because of particle size. More molecules will have more effect. Corrected volume V = V nb b is a constant that differs for each gas. Pressure Correction Because the molecules are attracted to each other, the pressure on the container will be less than ideal. Pressure depends on the number of molecules per liter. Since two molecules interact, the effect must be squared. P n V 2 observed P a ( ) Van der Waal s equation a and b are determined by experiment. a and b are different for each gas, bigger molecules have larger b a depends on both size and polarity.
Compressibility Factor The most useful way of displaying this new law for real molecules is to plot the compressibility factor, Z : For n = 1 Z = PV / RT www.tiwariacademy.in Ideal Gases have Z = 1 REAL GAS -- IDEAL GAS Very high temperatures and very low pressure help real gases simulate ideal gases. @ High Temperatures: Gas molecules move so quickly that there is not time to interact. @ Low Pressure: Gas molecules don t encounter each other very often. Standard Temperature & Pressure STP 1 atmosphere pressure (atm) 273 Kelvin (K) or 0 0 C Standard Molar Volume The volume occupied by one mole of gas at STP 22.4 L