Chapter 13 Liquids and solids

13.1 Intermolecular forces The physical properties of a substance generally depend on: the spacing between particles (atoms, molecules, ions) that make up the substance the forces of attraction among them.

Intermolecular forces determine a substances state at room temperature as well as melting point and boiling point

Reviewing what we know gases Solids Low density Highly compressible Expand to fill container High density Only slightly compressible Rigid (keeps its shape)

Intermolecular vs. intramolecular forces Inter- means between or among Intermolecular forces can hold together identical particles or two different types of particles Weaker than intramolecular forces (bonds)

Forces in water:

London Dispersion forces Occur in all substances Weak and short lived Only intermolecular force nonpolar molecules experience Atoms develop a temporary dipolar arrangement of charge as electrons move around the nucleus This instantaneous dipole can then induce a similar dipole in a neighboring atom.

Become stronger as the size of atoms or molecules increase

https://www.youtube.com/watch?v=1iYKajMsYPY

Dipole-Dipole Attraction Attraction between oppositely charged regions of polar molecules Polar molecule = Neighboring polar molecules orient themselves so that oppositely charged regions align

Hydrogen bonding Dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a fluorine, oxygen, or nitrogen atom

What types of intermolecular forces would be present in the liquid state of each of the following? CH3OH co2 h2o Kr NH3

Rank the following from lowest to highest boiling point CH3OH co2 h2o Kr NH3

13.2 Phase Changes Phase changes are physical changes no chemical bonds are broken/made

Energy requirements of phase changes Molar heat of fusion – energy required to melt 1 mole of a substance (water = 6.01 kJ/mol) Molar heat of vaporization – energy required to vaporize 1 mole of a substance (water = 40.7 kj/mol)

Liquid  steam: 40.7 kJ/mol Csteam = 1.70 J/goC CWater = 4.18 J/goC solid  liquid: 6.01 kJ/mol Cice= 2.10 J/goC

How much heat must be absorbed to melt 150.0 g of ice?

How much heat is released when 50.0 g of steam cools to 40oC?

How much energy is required to completely change 25 g of liquid water at 25 degrees Celsius to steam?

13.3 Liquids Kinetic molecular theory also applies to liquids and solids Must take intermolecular forces into account to apply it

Much denser than gasses Due to intermolecular forces holding particles together Fluidity – both gases and liquids are classified as fluids because they can flow and diffuse Liquids diffuse more slowly because intermolecular attractions interfere with the flow

Viscosity - measure of the resistance of a liquid to flow Attractive forces – stronger intermolecular forces = higher viscosity Particle size – larger molecules = higher viscosity Temperature – lower temperature = higher viscosity

Surface tension – the energy required to increase the surface area of a liquid by a given amount Caused by intermolecular forces pulling down on the particles on the surface of a liquid which stretches it tight like a drum

Cohesion – force of attraction between identical molecules Adhesion – force of attraction between molecules that are different

Evaporation To escape into the vapor phase a molecule must have enough energy to overcome the intermolecular forces of the liquid Rate increases as temperature increases Why does sweating cool you down?

Vapor pressure - the pressure exerted by a vapor over a liquid In a closed container will eventually reach equilibrium

https://www.youtube.com/watch?v=JsoawKguU 6A

Boiling point – temperature at which vapor pressure = atmospheric pressure

Predict which substance in each of the following pairs would have the largest vapor pressure H2O vs. CH3OH Ch3OH vs. Ch3Ch2Ch2oh

13.4 Solids Solid particles have as much kinetic energy as liquids or gasses but much stronger attractive forces between particles Limit the motion of particles to vibrations

Density of solids – almost always greater than density of liquids Exception = water

Crystalline solids – solid whose atoms, ions, or molecules are arranged in an orderly, geometric structure

Classify the type of solid:

Ionic solids Made of cation + anion Each ion is surrounded by ions of the opposite charge High melting point Brittle Held together by strong forces between ions

Molecular solids Fundamental particle is a molecule (relatively small) Melt at relatively low temperatures Held together by weak intermolecular forces

Atomic solid Fundamental particle is the atom Properties vary greatly Group 18 - low melting points Diamond - very high melting point

Metallic Bonding Metals are held together by nondirectional covalent bonds (called the electron sea model) among the closely packed atoms

Amorphous solids Particles are not arranged in a regular, repeating pattern Does not contain crystals Forms when molten material cools too quickly for crystals to form Glass Rubber Some plastics

Name the type of crystalline solid formed by each of the following: Ammonia Iron Cesium fluoride Argon sulfur

Which of the following is most likely to be solid at room temperature? Bao Hf O2