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Solids and Liquids

Solids and Liquids. Chapter 11. Summary. A substances state of matter depends on two things: The average kinetic energy of the particles (temperature) The strength of the intermolecular forces (IMFs) between the particles

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Solids and Liquids

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  1. Solids and Liquids Chapter 11

  2. Summary • A substances state of matter depends on two things: • The average kinetic energy of the particles (temperature) • The strength of the intermolecular forces (IMFs) between the particles • As temperature increases, the particles move farther apart (solid  liquid  gas) because the kinetic energy overcomes the IMFs holding the particles together

  3. Intermolecular forces • Defined as the attractive or repulsive forces between molecules, caused by partial charges • The type of bonding and structure of the substance are responsible for the type of intermolecular force involved and the strength of that force • The types of compounds are: • Nonpolar covalent • Polar covalent • Ionic **You will need to be able to draw a Lewis dot structure and determine whether a molecule is polar or nonpolar

  4. London Dispersion forces • This type of IMF is present in all substances but is only significant when the other types of IMFs are absent • This type of IMF is weaker than the others but is also dependent on the total number of electrons in a molecule (so may be strong enough for the substance to exist as a liquid or solid at room temperature) • More electrons = more polarizable, so the temporary dipole happens more often making the London force stronger than other types of IMFs in smaller molecules

  5. Dipole-Dipole forces • This type of force is only present in molecules that have a permanent dipole (polar covalent compounds) • This type of force is stronger than London forces in compounds of a similar size • This type of force is USUALLY responsible for a covalent compound being a liquid or solid at room temperature

  6. Hydrogen bonds • These are a special type of dipole-dipole force, found in molecules that contain H bonded to N, O, or F. • The H on one molecule is attracted to the N, O, or F of a neighboring molecule • Because the bonding within the molecule is extremely polar, hydrogen bonds between two molecules are generally stronger than normal dipole-dipole forces in compounds of a similar size

  7. Ion-Dipole forces • Ion-dipole forces are responsible for the solubility of ionic compounds: • If an ionic compound dissolves in water, the ion-dipole force (ion to water attraction) is stronger than the ionic bonds holding the crystal together • If an ionic compound is insoluble in water, the ionic bonds are stronger than the ion-dipole force that would pull the ions away from the crystal

  8. Ionic Bonds • Ionic Bonds are stronger than ANY intermolecular force, which is why ionic compounds are always solids at room temperature with extremely high melting points (and high boiling points).

  9. The Liquid State • In the body of a liquid, molecules are being pulled in all different directions by the intermolecular forces. At the surface of a liquid, molecules are only being pulled from the sides and below, thus creating surface tension (the tendency of a liquid to minimize it surface area). THE STRONGER THE INTERMOLECULAR FORCES IN A LIQUID, THE GREATER THE SURFACE TENSION. • Liquids also have a resistance to flowing, or viscosity. LIQUIDS WITH STRONG INTERMOLECULAR FORCES TEND TO HAVE A HIGHER VISCOSITY THAN THOSE WITH WEAKER INTERMOLECULAR FORCES. Viscosity is also temperature dependent. Higher Temperature = Lower viscosity • Heat of vaporization is the energy needed to convert the liquid to a gas. THE STRONGER THE INTERMOLECULAR FORCES, THE HIGHER THE HEAT OF VAPORIZATION. • Vapor pressure is the pressure exerted by the vapor form of a liquid over the surface of a liquid in a closed container. THE STRONGER THE INTERMOLECULAR FORCES, THE LOWER THE VAPOR PRESSURE • Boiling point is the temperature at which the liquid will boil (usually at standard pressure). THE STRONGER THE INTERMOLECULAR FORCES, THE HIGHER THE BOILING POINT.

  10. The Solid State • There are two main types of solids: • Amorphous – solids in which there is no extensive ordering of the particles • Crystalline – solids that display a highly ordered arrangement of particles in a 3D structure known as a crystal lattice (there are repeating units called unit cells) • Five types of crystalline solids are known • Atomic solids – individual atoms are held in place by London dispersion forces. The noble gases are the only known atomic solids to form • Molecular solids – lattices composed of molecules held in place by any type of IMF (water and sugar are examples) • Ionic solids – lattices composed of ions held together by attraction between oppositely charged ions (all ionic compounds exist as ionic solids) • Metallic solids – lattices formed by metallic bonding (electrons are delocalized and allowed to move throughout the entire sample (all metals exist as metallic solids) • Covalent network solids – lattices formed when covalent bonds form one giant molecule; these crystals are stronger than all molecular solids and some ionic solids (examples include diamond and SiO2)

  11. Heating curves • A heating curve is a graph of the temperature of a system versus the amount of heat added. See Figure 11.19 on p. 450 for a sample heating curve. • On each sloped segment of a heating curve, the sample of matter, in a set state of matter, is rising in temperature. The formula for calculating the change in enthalpy along a sloped segment is : ΔH = (m)(c)(ΔT) (where m = mass of the sample, c = specific heat capacity of the sample, and ΔT = the temperature change imposed on the sample) • One each flat segment of a heating curve, the sample of matter is undergoing a phase change. The lower flat segment corresponds to melting and the upper flat segment corresponds to boiling. The formula for calculating the change in enthalpy along a flat segment is: ΔH = (mol)(ΔHtrans) (where mol is the number of moles of the substance and ΔHtrans is the molar heat of fusion (if melting) or the molar heat of vaporization (if boiling) • Page 451 Sample Exercise 11.4 shows how to calculate ΔH for a set of segments on a heating curve.

  12. Phase Diagrams • A phase diagram is a graph representing a substance’s states of matter to temperature and pressure. The diagram allows us to predict which state of matter a substance will assume at a certain combination of temperature and pressure. • See Figures 11.27 and 11.28 on page 458 • You should be able to • Identify three general areas and label them as solid, liquid or gas • Identify the segment that corresponds to melting-freezing equilibrium • Identify the segment that corresponds to evaporation-condensation equilibrium • Identify the point that corresponds to equilibrium between all 3 states of matter and all 6 phase changes (the triple point) • Identify and define the critical point

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