Chem II Exam 1 Review PDF

# Chapter 11: Intermolecular Forces and Liquids 1. Intermolecular forces - forces that are responsible for the attraction of molecules to each other. - All kinds of intermolecular forces are called van der Waals Forces. - Covalent Bonding vs. Intermolecular Forces - Intermolecular f...

# Chapter 11: Intermolecular Forces and Liquids 1. Intermolecular forces - forces that are responsible for the attraction of molecules to each other. - All kinds of intermolecular forces are called van der Waals Forces. - Covalent Bonding vs. Intermolecular Forces - Intermolecular forces are weaker than the bonds that make compounds. - Covalent bonds are strong. 2. Weak Intermolecular Attraction - Polar - partial charges - Nonpolar - because bonds cancel out due to geometry - Polarity of individual bonds is due to electronegativity differences between bonds. 3. Dipole-dipole forces - bind molecules having permanent dipoles to one another. - Covalent bonds - $C = O$ - $C - O$ - Non-polar mainly have only carbons and hydrogens - Polar substances are likely to dissolve in polar solvents. - *Like dissolves like* - H-Bonding - When a H atom is covalently bonded to N, O, or F on a neighboring group, it is attracted to N, O, or F. - It is not a covalent bond. $H - C - O - H$ # Dispersion Forces - All molecules can do dispersion, but non-polar bonds can only do dispersion. - Polar can do dispersion and H-bonding. - Momentary dipole - the fluctuation of electrons. - The attraction is $O^+ \longrightarrow O^-$. - $H - C - H$ - $H - C - H$ - Dispersion are the weakest. - Dipole-dipole is stronger than dispersion forces - H-bord is the strongest. # Lecture 7 Outline - Intermolecular forces - attractions between forces as (the weaker force between covalent and ionic birds) - Van der Waals forces - weak forces due to temporary dipole. - All molecules have these, but this force is only present in nonpolar molecules. - Polar molecules are van der Waals and do H-bonding. - 3 areas where intermolecular forces exert a profound effect: - Melting point, boiling point, ad energy conversion to solid or liquid. - Solubility of gas, liquid, and solid. - Determining structure and behavior of important molecules like DNA and protein (hydro-carbon chains). - Causes of polarity of individual bonds. - Partial negative and positive charges. - Electronegativity - the ability of an atom to attract shared electrons to itself. Fluorine has the highest electronegativity. - Dipole-Dipole force - Forces between molecules with permanent dipole forces. - $H - Cl$ # Strengths of Intermolecular Forces - Weakest < Strongest - Van der Waals forces < dipole - dipole < hydrogen bonds < ion-dipole - The greater the forces of attraction between the molecules in a liquid, the greater the energy that must be supplied to separate them. Polar compounds have a higher value for their enthalpy of vaporization than nonpolar compounds with similar molar masses (polar-higher, nonpolar-less). - Boiling point - the greater the polarity, the higher the temperature required for the liquid to boil. # Chapter 11.6: Properties of Liquids ## a. Vaporization and Condensation (L->G) (G->L) - Molecules escape from the liquid surface and enter the gaseous state. - Breaking IM forces require energy. - Vaporization is an endothermic process. - Energy required to vaporize = molar enthalpy of vaporization, $\Delta H_{vap}$. - For condensation, things are equal, but opposite of vaporization. - Ex #1 - $KJ/14.6gC_4H_{10}/1molC_4H_{10} = -30.7kJ = -17.6 kJ$ - $78.0gC_4H_{10}/1molC_4H_{10}$ ## b. Vapor Pressure - The pressure above the liquid. - Equilibrium vapor pressure is the pressure exerted by a vapor over a liquid in a closed container. - Rate of evaporation = the rate of condensation. - Volatility - tendency of its molecules to escape from the liquid phase and enter the vapor phase at a given temperature. - The higher the equilibrium vapor pressure, the more volatile the substance. ## c. Boiling Point - Of a liquid is the temperature at which its vapor pressure is equal to the external atmospheric pressure. - Normal boiling point = 760 mmHg (1 atm) - $Vap. Press. (mmHg)$ - $0^oC$ - $100^oC Temp.$ # Viscosity - The measure of a fluid's (how thick or thin a fluid is) resistance to flow. - Surface tension - is the physical property of a liquid's surface to resist an external force, the stronger the IMF in a liquid, the higher the surface tension. - Weakest attraction have higher vapor pressure. - Boiling point, surface tension, and viscosity increase as intermolecular forces increase. - Higher viscosity, stronger intermolecular forces. # Lecture 3 Outline - Energy required to vaporize a sample = molar enthalpy of vaporization, $\Delta H_{vap}$. - Condensation is the process through which molecules change from gaseous phase into liquid (the opposite of vaporization). - Condensing produces a negative (exothermic process) $\Delta H_{cond}$. ## Equilibrium Vapor Pressure Change with Temperature - As the temperature increases, the equilibrium vapor pressure of liquid increases. - Pressure - $160$ - $100$ - $Temp. (°C)$ - Weaker forces = higher vapor pressure. ex. mosquito dancing on water. - Stronger forces = lower vapor pressure. ex. human finger in water. ## Relationship Between Intermolecular Forces and Viscosity - Higher viscosity, stronger molecular force. - Boiling point, surface tension, and viscosity increase as intermolecular forces increase. # Chapter 13: Solutions - A solution is a homogeneous mixture of two or more substances in a single phase (no pure substance). - The component present in largest amount is identified as the solvent, and the other component as the solute. - $Ex. H_2O$ solvent & Solute = solution. ('Substance being dissolved) - Colligative properties are properties of solutions that depend only on the number of the solute particles per solvent molecule and not to the identity of the solute. - Molality (m) of a solution is defined as the amount of solute per kilogram of solvent. - $m = moles$ of solute - $Kilograms$ of solvent - *Molarity is (M) mol solute/ L Solution* # Mastery Ex #1 - $grams$ $NH_4C_2H_3O_2$ $m = 0.618 = mol$ $NH_4C_2H_3O_2$ - Given: $375 grams$ - $Kilo: 1000$ - $375g$ $H_2O/1kg = 0.375kg$ - $1000g$ - $LS: S-TL$ - $mol$ $NH_4C_2H_3O_2 = (0.618)(0.375) = 0.232 mol$ - $g$ $NH_4C_2H_3O_2/0.232 mol | 77.0g$ $NH_4C_2H_3O_2 = 17.9g$ $NH_4C_2H_3O_2$ - $1 mol$ # Lecture 5: The Solution Process - Solubility - is the concentration of solute in a saturated solution in equilibrium with undissolved solute. - $Ex.$ ## Liquids Dissolving in Liquids - If two liquids can be mixed to form a homogeneous mixture, they are said to be miscible. In contrast, immiscible liquids do not fully mix; they exist in contact with each other as separate layers. - *Miscible = mixable, and remember "like dissolves like"* - Note: The more carbons that a molecule has, the more non-polar it is; also that "like dissolves like", and that water is polar. # Henry's Law - Henry's Law expresses the relationship between the solubility of the gas and the gas pressure (gas is directly proportional to partial pressure of the gas, $P$). - $S_g = k_H P_g$ - $S_g$ = solubility of gas (usually mol/kg) - $P_g$ = partial pressure of the gas over the solution. - $k_H$ = Henry's Law constant. - Solubility of a gas is defined as the concentration of dissolved gas in equilibrium with the substance in the gaseous state - Carbonated soft drinks illustrate how Henry's law works. # Mastery Ex - $S_g = k_H P_g$ - $\longrightarrow S_g = k_H$ - $K_H = 4.78 \times 10^{-3}$ $mol/L \cdot atm$ - Given: $S_g = 9.63 \times 10^2 g/L$