Chapter 7 - Intermolecular Forces

Chapter Seven: Intermolecular Forces The Three States of Matter existing as a gas, liquid, or solid depends on: • the balance between the kinetic energy of its particles • the strength of the interactions between the particles Intermolecular Forces, Boiling Point, and Melting Point intermolecular forces are the attractive forces that exist between molecules; in order of increasing strength, these are: • london dispersion forces • dipole-dipole interactions • hydrogen bonding strength of the intermolecular forces determines: • if the compound has a high or low melting/boiling point • if it is solid, liquid, gas at a given temperature London Dispersion Forces london dispersion forces are very weak interactions due to the momentary changes in electron density in a molecule • the change in electron density creates a temporary dipole • the weak interaction between these temporary dipoles constitutes london dispersion forces • all covalent compounds exhibit london dispersion forces • the larger the molecule, the larger the attractive force, and the stronger the intermolecular forces Dipole-Dipole Interactions dipole-dipole interactions are the attractive forces between the permanent dipoles of two polar molecules Hydrogen Bonding hydrogen bonding occurs when a hydrogen atom bonded to O, N, or F is electrostatically attracted to an O, N, or F atom in another molecule; hydrogen bonds are the strongest of the three types of intermolecular forces Boiling Point and Melting Point boiling point: temperature at which a liquid is converted to the gas phase melting point: temperature at which a solid is converted to the liquid phase the stronger the intermolecular forces, the higher the boiling and melting point Liquid State: Vapor Pressure evaporation: the conversion of liquids into the gas phase endothermic: absorbs heat from the surrounding condensation: the conversion of gases into the liquid phase exothermic: gives off heat to the surroundings vapor pressure is the pressure exerted by gas molecules in equilibrium with the liquid phase. vapor pressure increases with increasing temperature. the boiling point of a liquid is the temperature at which its vapor pressure = 760 mmHg the stronger the intermolecular forces, the lower the vapor pressure at a given temperature Viscosity and Surface Tension viscosity: a measure of a fluid’s resistance to flow freely • a viscous liquid feels “thick” • compounds with strong intermolecular forces tend to be more viscous • substances composed of large molecules tend to be more viscous • ex: olive oil is more viscous than water since it’s a larger polar molecule surface tension: a measure of the resistance of a liquid to spread out the stronger the they be jumpin! intermolecular forces, the stronger the surface molecules are pulled down toward the interior of a liquid and the higher the surface tension. water has a very high surface tension because of its strong intermolecular hydrogen bonding. when objects seem to float on the surface of water, they are being held up by the surface tension only. The Solid State solids can either be crystalline or amorphous. • a crystalline solid has a regular arrangement of particles - atoms, molecules, or ions - with a repeating structure • an amorphous solid has no regular arrangement of its closely packed particles there are four different types of crystalline solids - ionic, molecular, network, and metallic Crystalline Solids ionic solid: composed of oppositely charged ions (NaCl) molecular solid: composed of individual molecules arranged regularly (H2O) network solid: composed of a vast number of atoms covalently bonded together (SiO2) metallic solid: a lattice of metal cations surrounded by a cloud of electrons that move freely (Cu) Amorphous Solids amorphous solids have no regular arrangement of their particles. they can be formed when liquids cool too quickly for regular crystal formation. very large covalent molecules tend to form amorphous solids because they can become folded and intertwined. examples include rubber, glass, and plastic. Specific Heat specific heat: the amount of heat energy (cal or J) needed to raise the temperature of 1 g of a substance by 1 °C the larger the specific heat, the less its temperature will change when it absorbs a particular amount of heat energy. example: how many calories are needed to heat a pot of 1600 g of water from 25 °C to 100 °C? step one: identify the known quantities and the desired quantity known quantities: desired quantity: mass = 1600 g ? calories T1 = 25 °C T2 = 100 °C ΔT = T2 - T1 = 100 °C - 25 °C = 75 °C the specific heat of water is 1.00 cal/(g.°C) step two: write the equation the specific heat is a conversion factor that relates the heat absorbed to the temperature change (ΔT) and mass. heat = mass x ΔT x specific heat step three: solve the equation substitute the known quantities into the equation and solve for heat in calories. Energy and Phase Changes: Converting a Solid to a Liquid heat of fusion: the amount of energy needed to melt 1 gram of a substance Converting a Liquid to a Gas heat of vaporization: the amount of energy needed to vaporize 1 gram of a substance Heating and Cooling Curves a heating curve shows how a substance’s temperature changes as heat is added a cooling curve shows how a substance’s temperature changes as heat is removed