Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 107	Fundamentals of Chemistry	Fall 2008
Lecture Notes: 2 December	© R. Paselk 2005
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Liquid - Properties, cont.

• Viscosity
  • due to van der Waals forces and long molecules
  • due to strong H-bonding (ethylene glycol, HOCH₂CH₂OH has a viscosity much like high MW syrup).
• Surface Tension: due to attraction of molecules of liquid for each other (diagram).
  • meniscus
  • detergents/surfactants and wetting (water striders, ducks etc.)

Water: water is so ubiquitous, and has so many important and even special properties, that we will talk a bit more about it.

Water is a very unusual, even incredible substance whose amazing properties are often unappreciated because of its ubiquity. Water's special properties include:

• extremely high mp and bp (0 °C & 100 °C K, compare to methane, -183 °C & -161 °C, with a MW of 16 vs. water's 18, see figure below or Figure 13.3, p 381 in text.)

• a high heat capacity (18 cal/°C mol vs. 8 cal/°C mol for methane)
• a high viscosity
• its solid form is less dense than the liquid form at the same temperature (ice floats on water - very rare)
• it has a high dielectric constant (78.5 vs. 1.9 for hexane)
• it has hi surface tension
• it has high heats of fusion (melting)/freezing and vaporization/condensation.

The high mp, bp, and heat capacity all predict relatively strong bonding between water molecules, H-bonding. Note environmental consequences - Earth's weather is much more pleasant because it is moderated by water, especially along coasts. Ice floating prevents "solid" seas, definitely a downer in environmental terms.

Water of course is a covalent structure: H-O-H. But what gives it its special properties is the polarity of its O-H bonds and the resultant dipole moments of the bonds and the molecule itself.

The water molecule itself is bent, with an angle of 104.5° between the hydrogens (compare to 109.5° for sp³ tetrahedron) as seen in Figure 8.21, 8.20 on p 242 of your text. Because of the very strong dipole moments of these bonds and the very small size of the hydrogen substituents on water, a slight degree of orbital overlap occurs between adjacent water oxygens and hydrogens to give partial covalent bonds known a H-bonds (effectively, can only form with O, N, & F). Note that the partial covalent character means that they are directional!

Within solid bulk water (ice) every water molecule is bonded to 4 others, as in the ice structure seen in Figure 13.19 on p 400 [overhead 2.3, VV] In liquid water the molecules are still bonded to a large degree (the heat of fusion for ice is only 13% of the heat of vaporization for ice, thus most of the H-bonds must survive melting). Of course in liquid water the bonds are very unstable (average lifetime about 10 psec = 10^-11 sec), exchanging constantly to give a "flickering cluster" structure. The various properties of water arise from this structure. (Note hi bp & mp, heat cap., viscosity, and, less obviously, that ice floats. This is because the molecules are in an open lattice rather than close-packed. G&G note that close-packed molecules would only occupy about 57% of volume. This would lead one to expect that ice would float "high." It doesn't because most of the structure remains in the liquid phase at 0° C.)

Water is also an excellent solvent for polar substances since its dipolar structure enables it to insulate them from each other and it can make good dipole-dipole and dipole-charge bonds. Figure 2.4 on pg. 38 shows the hexavalent liganding of water to sodium and chloride ions to form hydration shells (For sodium ions, the waters in the inner hydration-shell exchange every 2-4 nsec.). Anything which can H-bond will also of course be quite soluble.

Solutions

Definitions:

• Solution - a homogeneous mixture of substances.
  • The major component is known as the solvent.
  • The substances dissolved in it (by definition minor components) are solutes.
• Solubility - a measure of how much solute will dissolve in a given solvent under given conditions (default = 20 °C, 1 atm).
  • Liquids which dissolve in each other are called miscible.
  • Liquids which do not dissolve in each other are considered immiscible.

Solubility:

• All gases are completely soluble in each other.
• Liquid solutions
  • "Like dissolves like."
  • Solvation - the process of surrounding a solute with solvent molecules).
    • Hydration (solvation with water).
  • Saturation - the maximum amount of solute which can be in solution in equilibria with its pure state.
    • Supersaturation - dissolving solute in excess of saturation - unstable to the addition of solute (carbonated beverages, honey, etc.)
  • Temperature effects:
    • Gases decrease in solubility with increasing temperature. (example of oxygen solubility)
    • Many solids increase in solubility with increasing temperature, but not always true.
  • Pressure effects:
    • Gases increase in solubility with increasing pressure (example of carbonated beverages)
    • Solubility of solids generally uneffected by pressure, since both liquids and solids essentially incompressible.

Colloids: defined by particle size = 1.0 nm< colloid < 100 nm (particles in solution are 0.1 - 1.0 nm in diameter, whereas particles > 100 nm dispersed in a fluid are considered to be in suspension.) Colloids generally do not settle out.

• Biomacromolecules are often colloidal (proteins, DNA & RNA)
• particles in inks are sometime colloidal.

Reaction Rates

Collision Theory: We assume that particles must collide in order to react. Thus a first understanding of reaction rates is based on understanding what influences the frequency of collisions.

• Can have different molecularites. This refers to the numbers of molecules (or particles - can also be atoms or ions) involved in the collision.
  • unimolecular - only one molecule involved (there is no collision)
  • bimolecular - 2 body collision
  • termolecular - 3 body collision (this is very rare, essentially never get more than three body collisions)
• Since particles must collide to react, the rate will depend on concentration. Thus
  • rate = k[A][B] for the reaction of A and B.
• But not all collisions will lead to reaction. At least two additional factors are important:
  • Activation energy, E_a. This is the energy needed to overcome repulsion, bond strength etc. At any given temperature the number of particles with a particular energy is given by N = e^-(E_a/RT). Note that
    • as E_a increases the fraction decreases.
    • as T increases, the fraction with sufficient energy for reaction (E>E_a) increases. Arrhenius Equation, k = Ae^-(E_a/RT). Can plot fraction of particles with KE vs. KE. [sketch or overhead]. A crude relation, which holds for many average reactions is the so called Q₁₀. This states that for every 10 °C increase in temperature we get a doubling of reaction rate.
  • Orientation of collision. The steric factor or probability factor (p) is the fraction of collisions which have orientations favorable for reaction. Can range from very small numbers (one atom or molecule must hit another at a very specific place) to about 1.

• Note that for kinetics the stoichiometry of the reaction is not necessarily related to the rate law, rather the rate law is related to a single slow step in the reaction process.

Transition States and Reaction Progress (Reaction Coordinate) Diagrams.

Reaction with a negative free energy (-DG) - products are favored:

.

How do we interpret this diagram?

• The x-axis can be thought of a a time axis for a single reaction, starting with the reacting molecules and ending with the product molecules.
  • Thus as the reaction progresses the energy in the particles goes up (they become less stable) as they are "pushed together" in a collision.
  • The upward slope represents the energy of repulsion and bond stretching.
  • The downward slope represents the formation of the new bonds and the separation of the products.
  • The top of the peak represents the "transition state" (TS) a high energy combination of atoms which can go to either reactants or products.
• The y-axis represents the energy of the species during the reaction. In these diagrams we are plotting the so-called free energy, DG= delta G. The free energy tells us how much energy is available to do work, and so is the most common energy measure used by chemists and biologists.
• The energy difference between the reactants and the TS is called the activation energy, E_act. This is the energy the reacting species must have (as kinetic energy etc.) in order to overcome the barriers to reaction (repulsion, bond energy). The greater the value of E_act the slower the reaction, because fewer molecules have enough energy to react. If E_act is very low (not much of a hill) then the reaction will proceed readily and rapidly.
• The energy difference between the reactants and the products (DG) tell us how far the reaction will go (that is, the fractions of reactants and products in equilibrium with each other). If the change is negative (products have less energy than reactants as seen on the diagram above), then products are favored (the reaction will go to mostly product. If, on the other hand, the products are at a higher energy then DG will be positive, and the reactants will be favored (the reaction will stay mostly reactants, and little product will be made - see the diagram below). If DG is zero, then the reaction mixture will end up with equal amounts of reactant and product,

  Reaction with a positive free energy (+DG) - reactants are favored:

  

  • Thus the size and sign of DG tells us how far a reaction will proceed, while the size of E_act tells us how fast the reaction will be. Note that reactions can be thermodynamically very favorable (go nearly completely to product and release lots of energy), but kinetically unfavorable (they react very slowly). We say the reactants are thermodynamically unstable, but kinetically stable.
• Catalysts make reactions go faster without affecting how far they go (DG is unchanged - the equilibrium is unchanged). They do this by
  • making it easier to achieve the activation energy (they reduce E_act), or
  • change the mechanism to one with one or more lower E_act's.

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