Chem 1412 - Chapter 11

Covalent Molecules

  • Also known as molecular molecules.
  • Composed of two nonmetals.

Ionic Compounds

  • Electron Exchange: Involves an exchange of electrons between atoms. Metals lose electrons, while nonmetals gain them.

  • Intermolecular Forces:

    • Positive and negative charges attract, leading to strong forces between molecules.
    • Ionic compounds have strong intermolecular forces due to the high attraction between positive and negative charges.

  • Boiling and Melting Points: Ionic compounds typically have very high boiling and melting points.

    • Transitioning between phases (e.g., solid to liquid) requires overcoming the strong attractions between molecules.

Covalent Molecules

  • Electron Sharing: Involves the sharing of electrons between atoms.

    • Example: Carbon (4 valence electrons) sharing with Hydrogen (1 valence electron) in CH4CH_4.
    • Lines in Lewis structures represent bonds between two shared atoms, with each line representing two electrons.

  • Equal vs. Unequal Sharing (Polarity):

    • Electronegativity: Some elements (F, O, N, Cl, Br) have high electronegativity, meaning they have a strong attraction for electrons.
    • Unequal sharing: When atoms with significantly different electronegativities form a bond, the electrons are unequally shared. The atom with higher electronegativity pulls the electrons closer, forming a partial negative charge (δ−\delta^-) and a partial positive charge (δ+\delta^+).
    • Dipole: The unequal sharing of electrons creates a dipole, represented by a dipole vector.
    • Polar Bond: Characterized by unequal sharing of electrons.
    • Nonpolar Bond: Characterized by equal sharing of electrons.

  • Polar vs. Nonpolar Molecules:

    • Polar molecule requirements:
      • Presence of one of the five highly electronegative elements (F, O, N, Cl, Br).
      • Asymmetrical structure; if the molecule is symmetrical, the dipoles cancel out, and the molecule is nonpolar.

  • Covalent Intermolecular Forces:

    • Dispersion Forces (London Dispersion Forces):
      • Present in all covalent molecules.
      • Arise from temporary, instantaneous dipoles due to fluctuations in electron distribution.
    • Dipole-Dipole Forces:
      • Occur in polar molecules.
      • Result from the attraction between the partial positive end of one molecule and the partial negative end of another.
    • Hydrogen Bonding:
      • A special type of dipole-dipole interaction.
      • Requires a hydrogen atom bonded to a highly electronegative atom (O, F, or N).
      • Stronger than typical dipole-dipole forces.

  • Relationship between Intermolecular Forces and Boiling Point:

    • Stronger intermolecular forces lead to higher boiling points.
    • Dispersion forces are the weakest, so nonpolar molecules tend to have low boiling points and exist as gases or liquids.

  • Molecular State:

    • Ionic compounds: Solids due to strong intermolecular forces.
    • Polar molecules: Often liquids due to dipole-dipole interactions.
    • Nonpolar molecules: Typically gases or liquids due to weak dispersion forces.

Lewis Structures and Polarity

  • Drawing Lewis Structures:

    1. Draw the skeletal structure.
    2. Count the total number of valence electrons.
    3. Subtract the number of bonding electrons.
    4. Place lone pairs to complete octets (8 electrons) for each atom, starting with terminal atoms.

  • Octet Exceptions:

    • Hydrogen only needs 2 electrons.
    • Boron only needs 6 electrons.

  • Determining Polarity from Lewis Structures:

    • Check for lone pairs on the central atom.
      • No lone pairs: Generally symmetrical and nonpolar (with some exceptions).
      • Lone pairs: Generally asymmetrical and polar.

Relative Boiling Points

  • Factors determining relative boiling points:
    1. Strength of intermolecular forces (hydrogen bonding > dipole-dipole > dispersion).
    2. Mass: Higher mass generally leads to higher boiling points.

Solutions

  • Mixing Salt in Water: Adding salt to water raises the boiling point.

  • Molecular Interactions:

    • The interactions formed when substances mix determine their miscibility.
    • The strength of these interactions relative to the interactions within the pure substances also plays a vital role.

  • "Like Dissolves Like":

    • Polar solutes dissolve in polar solvents.
    • Nonpolar solutes dissolve in nonpolar solvents.

  • Nonpolar molecules: Held together by dispersion forces only, which are the weakest intermolecular forces.

  • Polar molecules: A molecule needs to be Asymmetrical to be polar.

  • Hydrogen Bonding with water: For hydrogen bonding, in addition to being polar, we need one of these three bonds, OH, FH, or NH. The hydrogen bonds between nitrogens are what sticks our two strands of DNA together.

Solutions and Intermolecular Forces

  • Boiling point increases when we mix something in with water and create a solution.

  • To start you need to determine, first of all, if it's polar or nonpolar.

  • Strength of intermolecular forces:

    • the stronger intermolecular forces, the higher the volume.

  • Metals to the left, nonmetals to the right. Metal + olefins (boron, silicon, germanium, arsenic, antimony, metaloids, right, between metals and nonmetals.

Solutions

  • When you put salt into water, the boiling point of the water goes up. So if regular water is 100 and you put the salt in and listen. Hey.

  • When you mix something in with water and create a solution, the boiling point should go up.

  • Electrolytes: The dissociation of ions forming multiple particles.

  • Nonelectrolyte: Polar molecules dissolved in water.

    • Electrolytes are ionic dissolved in water, and nonelectrolytes are polar molecules dissolved in water.

  • Solubility table: Top half soluble ionic compounds. Right. So we're determining which ionic compound dissolved in water, I.E., soluble and which do not have it in soluble. Yeah. So if the anion is nitrate, right, they're always soluble. If the anion is chloride, for example, chlorides are always soluble unless the cation is one of these three. So silver chloride, for example, will be insoluble.

Formulas:

  • Delta T and B: Boiling point, is equal to the Van Hoff factor times the boiling point constant.
  • Formula: delta t for freezing, again, the Van Haul factor times the freezing point constant times the. Delta t=i∗kf∗mt = i * k_f * m where:
    • ii is the van't Hoff factor,
    • kfk_f is the cryoscopic constant (freezing point depression constant), and
    • mm is the molality of the solution.

Terms:

  • Solvent - water.
  • Solute - a particle inside the solvent (ex: NaCl).
  • Solution - The product created. All wait, this one's solvent water. So is it polar or nonpolar.

Solubility

  • If I increase the temperature for solids, solubility is also increased.

  • If I increase the pressure, solubility increases, decrease the pressure, the solubility decreases. Okay okay yeah for solids and for gases. Okay What we're about to talk about relative to this. Information is something called colligative properties and those are what that basically means is that it's dependent on the number of solute particles in solution

Concentration:

  • When doing concentration make sure not to confuse the solute from the solution.
    Here's the information we're getting about the solvent. There's 750 mils of it volume and mass. So the density is 1.3. And we have the volume here.

  • Molarity=moles of soluteliters of solutionMolarity = \frac{moles \, of \, solute}{liters \, of \, solution}

  • Molality=moles of solutekilograms of solventMolality = \frac{moles \, of \, solute}{kilograms \, of \, solvent}

  • Weight percent=mass of solutetotal mass of solutionWeight \, percent = \frac{mass \, of \, solute}{total \, mass \, of \, solution}

  • Mole fraction=moles of solutetotal molesMole \, fraction = \frac{moles \, of \, solute}{total \, moles}

  • There are no pressure effects for for solids on you. Right. We're only gonna talk about temperature effects. So for gases, the pressure effect. Right.