AAC Chemistry Test Review - Polarity, IMF & Solubility PDF

# AAC CHEMISTRY Test Review: Polarity, IMF & Solubility ## Lewis Structures Lewis structures are depictions of molecules that show valence electrons as dots. Shared pairs of electrons (i.e. bonded) are drawn between the atoms sharing them. Unshared or lone pairs of electrons are represented by dots...

# AAC CHEMISTRY Test Review: Polarity, IMF & Solubility ## Lewis Structures Lewis structures are depictions of molecules that show valence electrons as dots. Shared pairs of electrons (i.e. bonded) are drawn between the atoms sharing them. Unshared or lone pairs of electrons are represented by dots located on one atom only. ### Examples 1. Draw the Lewis structure for methane, CH4. - N = - A = - S = 2. Draw the Lewis structure for phosphorus trichloride, PCI3. - N = - S = 3. Draw the Lewis structure for the carbonate ion. - N = - A = - S = ## Polarity of Bonds - The bonding pairs of electrons in covalent bonds are located between the pairs of the atoms sharing the electrons. - When the pull of each nucleus for the electrons is equally strong, the electrons are shared equally and are located on average halfway between the two nuclei. Recall that this type of bond is a nonpolar covalent bond. - When the pull of one nucleus for the electrons is stronger than the other, the electrons spend more time closer to the more electronegative nucleus. Recall that this type of bond is a polar covalent bond. - The more electronegative atom acquires a partial negative charge. The less electronegative atom therefore acquires a partial positive charge. - These partial charges are indicated by the following symbols: - δ - for the more electronegative element. - δ + for the less electronegative element. ### Examples - Ex: H2O The water molecule contains two dipole moments, sometimes just called dipoles. The dipole moments can be drawn in with arrow notation: 4. Determine bond polarity for each of the following bonds and draw the arrow notation for the dipole moment. - a. H-H - b. C-O - c. H-Cl - d. N-F - e. P-I 5. Draw the Lewis Structure for each molecule, draw the dipole moments for each polar bond. - a. H2S - b. SO2 - c. SiCl4 - d. CH3Cl - e. HF - f. PI3 ## Polarity of Molecules - Note that just because a molecule contains polar bonds, it may or may not be classified as a polar molecule. - The polarity of a molecule will depend on: - (1) existence of polar bonds - (2) shape of the molecule - (3) orientation of the polar bonds - (if none of the bonds are polar, the molecule is nonpolar*) ### Examples - Note: to help determine if the molecule will be polar, look for lone pairs on the central atom. If lone pairs are present, the molecule will be polar. - Ex: NH3 - The ammonia molecule contains three dipole moments. The NET dipole is the overall dipole moment of a molecule, resulting from the vector sum of all individual bond dipoles within the molecule. This arrow is drawn toward Nitrogen. 6. Draw the Lewis Structure for each of the following molecules, draw the dipole moment for each polar bond AND the NET dipole for the molecule - a. HCl - b. CO2 - c. CF4 - d. NH3 - e. CH2F2 - f. CH3N ## Intermolecular Forces (IMF) - **London Dispersion Forces (LDFs):** Temporary dipoles created by random electron movement in atoms/molecules. - Weakest intermolecular force. - Key Features: - Present in all molecules (polar and nonpolar). - Strength increases with molecular size and surface area. - **Dipole-Dipole Interactions:** Attraction between permanent dipoles in polar molecules. - Moderate strength. - Key Features: - Occur only in polar molecules. - Examples: HCl, CH3Cl. - **Hydrogen Bonding:** Attraction between a hydrogen atom bonded to a highly electronegative atom (N, O, F) and a lone pair of electrons on another molecule. - Strongest intermolecular force (but weaker than covalent/ionic bonds). - Key Features: - Responsible for unique properties of water (high boiling point, surface tension). - Found in molecules with H bonded to N, O, or F. - Examples: H2O, NH3, HF ## Practice Problems 7. Determine if the molecule is polar or non-polar and state the strongest IMF present in each of the following: - a. NF3 - b. SiH4 - c. CO2 - d. N2 - e. CBr4 - f. CH3OH - g. H2CO 8. Rank the following in terms of increasing boiling point: - LiCl - CH3OH - CH4 - O2 ## Solutions and Solubility - **Solution Definition:** A homogeneous mixture of a solute (substance being dissolved) and a solvent (substance doing the dissolving). - **Ionic Compounds in Water:** - Ionic compounds are of positive (cations) and negative (anions) ions held together by ionic bonds. - Dissociate into individual ions when dissolved in water (e.g., NaCl → Na+ + Cl-). - Not all ionic compounds are soluble. We use a solubility chart to determine which compounds are soluble. 9. Determine if each of the following compounds would be soluble or insoluble. - a. AgNO3 - b. LiBr - c. Ba(OH)2 - d. PbI2 - e. CaCO3 - f. Na2CO3 - g. Lead (II) chloride - h. Sodium sulfate - i. Magnesium sulfide - j. Barium chlorate - k. Copper (II) phosphate - l. Ammonium dichromate ## Precipitation Reactions - Precipitation reactions are a type of double replacement reaction. Two solutions of ionic compounds are mixed. One of the products of the reaction is an insoluble salt called a precipitate. Solubility rules are used to determine the identity of a precipitate. - ***Use the STAAR Solubility of Common Ionic Compounds in Water Chart on the back of your Periodic Table.*** - Write the products for these double replacement reactions and balance. Then, using the solubility rules above, identify the precipitate by a subscript (s) for solid (insoluble in water). All the other compounds should be marked with a subscript (aq) for aqueous (soluble in water). - Example: AgNO3 (aq) + LiCl (aq) → AgCl (s) + LiNO3 (aq) - 10. Predict the products of the following reactions. If a solid precipitate is formed, circle yes. CIRCLE the PRECIPITATE. - a. Na2CO3 + Ca(NO3)2 → - b. Pb(C2H3O2)2 + Na2S → - c. K3PO4 + MgCl2 → - d. (NH4)2SO4 + Bal2 →