CHEM 1212 1nd Edition Lecture 1Outline of Last Lecture Outline of Current Lecture I. States of Matter and Intermolecular ForcesII. Interactions between Ions and Molecules with a Permanent DipoleIII. Interactions between Molecules with a Dipolea. Dipole-Dipole Forcesb. Hydrogen Bondingc. Hydrogen Bonding and the Unusual Properties of WaterIV. Intermolecular Forces Involving Nonpolar Moleculesa. Dipole-Induced Dipole Forcesb. London Dispersion Forces: Induced Dipole-Induced DipoleV. A Summary of van der Waals Intermolecular ForcesCurrent LectureChapter 12.1-12.5I. States of Matter and Intermolecular Forcesa. Kinetic-Molecular Theory of Gas – assumes that gas molecules are widely separated and these particles can be considered to be independent of one anotherb. In ideal conditions, the properties of gas fall under the ideal gas law (PV=nRT)i. P = pressure of gasii. V = volume of gasiii. n = amount of gas (moles)These notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.iv. R = gas constant v. T = gas temperaturec. In real gasses, intermolecular forces (IMFs) are at work; if these forces are strong enough, they cause the gas to condense to a liquid and eventually a solidd. There is a large increase in volume when converting liquid to a gas, but no dramatic change in volume occurs when converting a liquid to a solidi. It can therefore be assumed that the molecules in liquid are about as closely packed as those in solidse. How do IMFs influence chemistry?i. Dipole-dipole forces: between molecules with permanent dipoles (molecule with permanent partial charges)ii. Dipole-induced dipole forces: polar molecules and nonpolar onesiii. Induced dipole-induced dipole forces: nonpolar molecules (also called London forces)II. Interactions between Ions and Molecules with a Permanent Dipolea. Ion-dipole forces: attraction between a positive or negative ion and a polar molecule b. Ion-ion forces are the most strongc. Coulomb’s Law: (a way to evaluate ion-dipole attractions) states that the force of attraction between two charged objects depends on the product of their charges dividedby the square of the distance between themd. Example of the interaction between an ion and a polar molecule:i. Hydration of ions (enthalpy of solvation/enthalpy of hydration)ii. Has a substantial enthalpy change, but it cannot be measured directlyiii. The enthalpy of hydration depends on the charge of the ion and 1/d (d = distance between center of ion and oppositely charged pole of dipole)III. Interactions between Molecules with a DipoleDipole-Dipole Forcesa. Dipole-Dipole Interaction: when one polar molecule interacts with another dipole molecule (either different of the same kind) the positive end of one molecule will be attracted to the negative end of the other moleculei. These attractions influence the evaporation of a liquid and the condensation of gas (each requires an energy change1. Evaporation requires the input of energy (enthalpy of vaporization) this change has a positive enthalpy; so this process is endothermic2. The enthalpy change for condensation is the opposite (exothermic) so it has a negative enthalpyb. The greater the force of attraction between the molecules, the greater the energy that must be applied to separate them i. So we would expect polar compounds a higher value of enthalpy of vaporization than nonpolar compounds with similar molar massesc. Boiling Pointi. As the temperature of a liquid is raised, its molecules gain kinetic energyii. Once enough kinetic energy is attained (at the liquid’s boiling point) the molecules have sufficient kinetic energy to break the IMF’s between themselves and neighboring moleculesiii. For molecules of similar molar mass, the greater the polarity of the molecule, the higher the boiling pointd. Solubilityi. “Like-dissolves-like” meaning polar molecules are likely to dissolve in polar solvents, and nonpolar molecules are likely to dissolve in nonpolar solventsii. Water is a polar solventHydrogen Bondinge. In general, the boiling point of hydrogen compounds (i.e. CH4, SiH4…etc.) increase with increasing molar massf. There are a few exceptions to this rule:i. Hydrogen compounds containing Nitrogen (N), Fluorine (F), or Oxygen (O), have much greater electro-negativities and therefore a much higher boiling point that would be assumed from their molar masses ii. The high electronegativity of these compounds is due to the large difference in electronegativity between H (2.2) and N (3.0), O (3.5), or F (4.0)g. In these cases, the more electronegative atom (N,O, or F) takes on a strong negative charge and the hydrogen atom acquires a strong positive chargeh. A hydrogen bond is the unusually strong attraction between an electronegative atom with a lone pair and the hydrogen atom of the N---H, O---H, or F---H bondi. The hydrogen in this molecule is then available to make a bridge between the N, O, or F in its molecule and the N, O, or F of another moleculei. Examples:1. N---H - - - :N--- F---H - - - :N---2. N---H - - - :O--- F---H - - - :O---3. N---H - - - :F--- F---H - - - :F---4. O---H - - - :N---5. O---H - - - :O---6. O---H - - - :F---*** The dotted line after the hydrogen represents the hydrogen bond***Hydrogen Bonding and Unusual Properties of Waterj. Water’s unique properties are due to its strong hydrogen bonding; each water molecule can participate in four hydrogen bondsk. Ice is an example where water molecules bond together through hydrogen bonds and this creates a tetrahedral structure; as a result is made up of tons of little cages of water molecules with lots of free space available l. The free space in the ice is the reason that it 10% less dense than water, which explains why ice floats on waterm. When ice melts, the hydrogen bonds break down causing the increased density of watern. Water’s density reaches a maximum at about 4 degrees Celsius and then declines in density with increasing temperature like other liquidso. This is why lakes/bodies of water freeze from the surface downi. Once the whole body of water reaches 4 degrees Celsius (all at same density), some water begins to freeze and it floats to the top due to its greater densityp. Hydrogen bonding is also the reason for waters high boiling point (High specific heat)i. It takes a greater input of energy (heat) to break the hydrogen bondsii. This is the reason that bodies of