CHEM 16.1 REVIEWER PDF

CHEM 16.1 REVIEWER A. Experiment #6: Calorimetry B. Experiment #7: Paper Chromatography C. Experiment #8: Dynamic Equilibrium and Heats of Solution D. Experiment #9: Distillation of Rubbing Alcohol E. Experiment #10: Gases F. Experiment #11: Colligative Properties G....

CHEM 16.1 REVIEWER A. Experiment #6: Calorimetry B. Experiment #7: Paper Chromatography C. Experiment #8: Dynamic Equilibrium and Heats of Solution D. Experiment #9: Distillation of Rubbing Alcohol E. Experiment #10: Gases F. Experiment #11: Colligative Properties G. Experiment #12: pH, Conductivity, and Relative Strengths of Acids and Bases Calorimetry process of measuring the amount of heat released or absorbed during a chemical reaction 1st law of thermodynamics: the total internal energy of the universe is constant = energy can neither be created nor destroyed, it can only be converted from one form to another o Law of Conservation of Energy = universe is an isolated system Types of Calorimetry Constant Pressure Calorimetry = Coffee Cup Calorimeter or Styroball Calorimeter Constant Volume Calorimetry = Bomb Calorimeter Calorimetry set-up HCl + NaOH ® NaCl + H2O Adiabatic system: no transfer of heat from the system to the surroundings Placed an insulator = Styrofoam ball o No easy absorbance/release of heat from the system o Very low heat capacity § Heat capacity = tendency to absorb/release heat qisolated system = qrxn + qcal qrxn = (∆Hrxn)(nLR) qrxn + qcal = 0 qcal = C∆T ; where C = mc or nCm Thus, qrxn = -qcal Thus, (∆Hrxn)(nLR) = -Ccal∆T If (+) ∆H à absorbed heat à Endothermic; temperature decreases If (-) ∆H à gives off heat à Exothermic; temperature increases Procedure Calibration (Determining heat capacity of calorimeter) Ccal is highly dependent on the amount of matter Need to calibrate because Ccal is not constant in all styrofoams; and in test tubes ∆Hrxn of HCl + NaOH ® NaCl + H2O = -55.85 kJ/mol "∆#$%& ∙()* o Ccal = ∆+ Determination of ∆Hrxn "/012⋅∆+ ∆𝐻-.( = ( )* HCl + NaOH ® NaCl + H2O and HNO3 + NaOH ® NaNO3 + H2O should have the same ∆Hrxn given that they both have the same NIE (H+ + OH- ® H2O) Troubleshooting There should be no space in between the test tube and the Styrofoam so the heat will not escape Don’t change test tubes/styrofoams in the course of the experiment because the amount of matter calibration will change as well Paper Chromatography Chromatography “chroma” + “graphein” = to illustrate/write color A technique for separation of mixture Developed by Mikhail Tsvet o In his experiment, he added plant extracts (with color) on top of calcium carbonate powder. Then he added water to it and after all water has passed on the packed powder, separation of color is formed o Each color corresponds to a certain compound Types of Chromatography o Column Chromatography and Planar Chromatography Chromatographic Set-up Mobile Phase o Moving part (H2O) Stationary phase o Staying part (paper) Sample mixture o To be separated How does separation occur? Two competing factors: There is a competition between the attraction of the sample to the solvent and to the paper. o Solubility to solvent (mobile phase) o Adsorption to paper (stationary phase) For paper chromatography o Higher solubility to solvent: moving faster o Higher adsorption to stationary phase: moving slowly Primary reason of solubility and adsorption: intermolecular forces of attraction (IMFA)  Separation is based on polarity o Molecular polarity § Due to tendencies of some element to attract more electron (electronegativity), they tend to disrupt electron cloud molecule, which leads to partial charge formation o Polarity is dependent on: § Electronegativity of atoms § Geometry o General rule of polarity: Like dissolves like o Derivatives of Polarity: IMFA (arranged according to decreasing strength) § Ion-ion § Ion-dipole § H-bonding § Dipole-dipole § Ion-induced dipole § Dipole-induced dipole § London dispersion forces Intermolecular Forces of Attraction London Dispersion Force o For all types of molecules Dipole-induced dipole and Ion-induced dipole Dipole-dipole o Polar + polar Hydrogen bonding o Special type of dipole-dipole (H-O, H-N, H-F)  Ion-dipole o Ion + polar molecules Ion-ion o Ion + ion (ionic) Types of Chromatography based on polarity of stationary phase Normal phase chromatography o Stationary phase is polar Reverse phase chromatography o Stationary phase is nonpolar Experiment made use of a Planar Normal Phase Chromatography (paper chromatography) o Stationary phase: water adsorbed in cellulose (cellulose is in paper) o Mobile phase: 1% NaCl solution in water Paper Chromatography Set-up  Chromatographic plate (chromatogram): filter paper with dye spots o Where the sample mixture to be separated is spotted o Stationary phase which is the filter paper Developing solvent: 1% NaCl solution o The mobile phase (1% NaCl solution) o Moves up the stationary phase via capillary actions Saturation Filter paper o Paper added besides the chromatographic plate o Used to ensure the system inside the chromatogram is saturated (prevent volatilization) Chromatographic Chamber: lid + beaker o Beaker + Lid = Where the chromatographic run is contained o Must be closed by lid to ensure no solvent would be volatilized Retention Factor Degree of attraction of components to the mobile phase If the component is: o More polar = lower Rf o More nonpolar = higher Rf Rf values are qualitative (used for comparison) Based on Rf Values: Allure Red < Sunset Yellow -?4@9:;] = 4: 4: Rate of reactions Indicates how fast a reaction occur Given by rate of change of concentration over time (derivatives) Le Chatelier’s Principle If a dynamic equilibrium is disturbed, the system would react with the change to attain equilibrium again (counteract with the change) Aims to restore balance Proposed by Henry-Louis Le Chatelier Two types of equilibria Acid-base equilibria o Effect of pH § ­pH = ¯H+ § ¯pH = ­H+ o Color changes are observed o We collected colored pigments (dyes) which contain anthocyanins that serve as natural indicators (have similar properties as phenolphthalein) Solubility equilibria o Effect of the nature of solute and solvent (miscible and immiscible) § Solubility occurs when the IMFA of the solute-solvent > solute-solute & solvent- solvent IMFA § Ex. Water & oil Water-water’s dominant IMFA is H-bonding Oil-oil’s dominant IMFA is LDF Water-oil’s dominant IMFA is dipole-induced dipole H-bonding > Dipole-induced dipole > LDF Thus, they would not interact with each other and form an immiscible liquid § Ex. Water & alcohol Water-water’s dominant IMFA is H-bonding Alcohol-alcohol’s dominant IMFA is H-bonding Water-alcohol’s dominant IMFA is H-bonding Thus, they would interact with each other and form a miscible liquid o Change in strength of the IMFA leads to change in the volume of solvent mixture o System wants to attain highest IMFA possible because it is most stable, and its harder to break Acid-base equilibria example Reaction: H+ + In ßà HIn+ ­pH = ¯H+ = reverse reaction ¯pH = ­H+ = forward reaction 𝑝𝐻 = −log [𝐻G ] When added HCl = forward reaction = ¯pH H+ + In à HIn+ ln = color A HIn+ = color B When added NaOH = reverse reaction = ­pH H+ + OH- + In à HIn+ H+ + OH- would form water and thus, would lessen H+ HIn+ à H+ + In Solubility Equilibria Effect of Nature of Solvent and Solute Solute and Solvent Glycerol: (polar; H-bonding) Naphthalene: (non-polar; LDF) Sodium Chloride: (ionic; ion-ion) Water: (polar; H-bonding) Ethanol: (polar; H-bonding) Methanol: (polar; H-bonding) Methanol is more polar than ethanol because the presence of C atoms also affects the polarity of the molecule. Less C = more polar; More C = more nonpolar Toluene: Thus, Glycerol + Water = Miscible (both polar) Glycerol + Naphthalene = Insoluble (polar and nonpolar) Ethanol + Naphthalene = Slightly Soluble (even if ethanol is polar, it still has an ethyl (C2H5) group that is nonpolar, thus attracting nonpolar solutes) Effect of Solubility examples I2 + H2O will not dissolve completely because I – I interaction is stronger (ion-ion) but will still dissolve because of I + I Û I2 I2 + KI in H2O will dissolve because KI + I2 Û I3- (represented by the change in color) Effect of temperature examples NH4Cl + water = endothermic (lumamig) NH4Cl + ∆H Û NH4+ + Cl- ­ temperature = more dissolved = forward reaction Ca(OH)2 in H2O = exothermic (uminit) Ca(OH)2 Û Ca2+ + 2OH- + ∆H ¯ temperature = more dissolved = forward reaction Changes in Volume and Heat H2O + MeOH (methanol) ∆𝐻;?H@:I?( = J ∆𝐻;?H@:7";?H@:7 + J ∆𝐻;?HL7(:";?H@:7 + J ∆𝐻;?HL7(:";?HL7(: ∆𝐻;?H@:I?( = J[+𝐵𝐵] + J[−𝐵𝐹] + J[+𝐵𝐵] IMFAsolute-solvent > IMFAsolute-solute + IMFAsolvent-solvent Thus, it is an exothermic reaction, ­T, ¯V Acetone + EtOH Acetone IMFAsolute-solvent < IMFAsolute-solute + IMFAsolvent-solvent Thus, it is an endothermic reaction, ¯T, ­V because the substances do not want to interact Response to IMFA ­IMFA = ¯V ¯IMFA = ­V Response to Concentration A + B ßà C + D More A: forward reaction o To reduce A, increase C More C: reverse reaction o To reduce C, increase A Response to Temperature Exothermic: A + B à C + heat o Decrease temperature to form heat Endothermic: A + B + heat à C o Increase temperature to absorb heat, to compensate with high temperature Response to Pressure ¯V, ­P = forward reaction (less gass molecules) ­V, ¯P = reverse reaction (more gas molecules) Distillation of Rubbing Alcohol Distillation Separation of miscible solutions o Miscible: unsure of the components Separation can be done by using the differences in the boiling points of the components Types of Distillation Set-up Simple distillation o used for separating liquids boiling below 150 °C at 1 atm from either nonvolatile impurities or another liquid with boiling point that is at least 25 °C from the first Vacuum distillation o used for separating liquids boiling above 150 °C at 1 atm from either non-volatile impurities or another liquid with boiling point that is at least 25 °C from the first Fractional distillation o used for separating liquids whose boiling points differ by less than 25 °C Steam distillation o used for separating liquids that are insoluble or slightly soluble in water Simple Distillation set-up Distilling flask: contains the mixture or solution How does Distillation occur? Due to difference of IMFA of the mixture’s components o For instance, in Rubbing alcohol, it is composed of H2O and Ethanol/Isopropyl Alcohol o However, Ethanol/isopropyl alcohol has weaker IMFA compared to water o Therefore, easier energy is needed to induce evaporation to it o Thus, it evaporates at lower temp compared to water Important Notes Temperature should not be greater than 100 C because by then, water would also evaporate You remove the first mL in the distillate because you assume they’re impurities There’s a hole in the condenser for air circulation, to avoid too high pressure that would lead to the breakage of the glassware o Because of this, there is an assumed loss of heat (energy is not fully conserved) Gases have weak IMFA, they are loosely connected no definite shape/volume; compressible have their own o pressure o volume o temperature o moles Gas Law Formula Description Boyle’s Law 𝑃P 𝑉P = 𝑃R 𝑉R At constant T, as ­P, ¯V Charles’ Law 𝑉P 𝑉R At constant P, as ­V, ­T = 𝑇P 𝑇R Gay-Lussac’s Law 𝑃P 𝑃R At constant V, as ­P, ­T = 𝑇P 𝑇R Combined Law 𝑃P 𝑉P 𝑃R 𝑉R Obtained by combining Boyle’s, = 𝑇P 𝑇R Charles’, and Gay-Lussac’s Law Ideal Gas Law 𝑃𝑉 = 𝑛𝑅𝑇 𝐿𝑎𝑡𝑚 𝑅 = 0.0821 𝑚𝑜𝑙𝐾 Objectives of the experiment The presence of gas law: the effect of presence of moles of gas molecules to pressure and volume and response to temperature How to use a eudiometer o Eudiometer is a classical instrument used to collect gas particles from a reaction to determine the volume of the gas Eudiometer Same as a barometer Solvent: water Three cases = three different formulas ℎ 𝐶𝑎𝑠𝑒 𝐴 = 𝑃:-8ff74 8I- = 𝑃g8- − 13.6  𝐶𝑎𝑠𝑒 𝐵 = 𝑃:-8ff74 8I- = 𝑃g8- ℎ 𝐶𝑎𝑠𝑒 𝐶 = 𝑃:-8ff74 8I- = 𝑃g8- + 13.6 * The discrepancy in level of water (represented by h) is due to difference in pressure of the outside (atmospheric pressure) and inside (pressure of trapped air) Determining pressure in eudiometer In the experiment, we were interested in the pressure of captured H2 gas in the reaction of Mg and HCl Mg + 2HCl à H2 + MgCl2 We are interested in the pressure of H2 only but water has a contribution to pressure due to vapor pressure Therefore, 𝑃4-k 8I- = 𝑃:-8ff74 8I- − 𝑃l8:7- L8f?- Determining Moles (n) of H2 gas PV = nRT Colligative Properties Colligative properties Properties of solutions which are dependent on the concentration of solute molecules/ions independent from identity of solute Types of colligative properties Boiling Point Elevation (BPE) Freezing Point Depression (FPD) Vapor Pressure Lowering (VPL) Osmotic Pressure (OP) Relationship of BPE and FPD to Concentration ∆𝑇 = 𝑖𝐾𝑚 ∆𝑇g = 𝑖𝐾g 𝑚 ∆𝑇n = −𝑖𝐾n 𝑚 Where: T = change in temperature compared to original solution K = proportionality constant (different for BP and FP) o Kf of water= 1.86 C/m o Kb of water = 0.512 C/m o Dependent to solvent, not solute o Determined empirically m = molality o?H7; ;?H@:7 o 𝑚 = pq ;?HL7(: i = Van’t Hoff factor o Ratio between the actual concentration of particles produced when the substance is dissolved and the concentration of a substance as calculated from its mass o Van’t Hoff Factor is dependent on the type of electrolyte the solute is § For Strong Electrolytes, i = total ions it produces upon dissolution § For Weak Electrolytes, i = less than the total ions it produces upon dissolution  § For Non-electrolytes, i = 1 § Example: (NH4)3PO4, i = 4 (-) sign for FPD designates its decrease Thus: Boiling Point Elevation o ∆T is (+) “elevation” à harder to boil Freezing Point Depression o ∆T is (-) “depression” à harder to freeze Determination of i ∆+ 𝑖 = ∆+ rstuvws& &s&xtxyv$stzvx 𝑖 = 2 theoretically in the experiment 𝑖 > 2, due to instrumental errors 𝑖 < 2, due to high molality concentration; oversaturated solutions (this is why we’d rather have diluted/unsaturated solutions) Addition of salt to the ice bath Salt decreases the freezing point of ice Induces melting of ice Melting of ice is endothermic More ice melting = more endothermic reaction Would lower the temperature of the solution Lowered temperature = achieve freezing point pH, Conductivity, and Relative Strengths of Acids and Bases Definitions of Acids and Bases Arrhenius Definition o Acid: gives off H+ § Ex: HCl, HNO3 o Base: gives off OH- § Ex: NaOH, NH4OH o However, there are compounds that are similar to properties of acids and bases that do not have OH- Bronsted Lowry Definition o Acid: proton donor § Ex: HCl, HNO3 o Base: proton acceptor § Ex: NaOH, NH3 o However, these definitions cannot explain acids and bases in inorganic compounds Lewis Definition o Acid: electron acceptor o Base: electron donor  Properties of acids and bases Conductivity Test Acids and bases have electrolytic properties = conduct electricity Strong acid/base = strong electrolytes = all ions dissolve Weak acid/base = weak electrolyte = partial dissociation o Glacial acetic acid does not conduct electricity alone, but when diluted, it conducts electricity H+ and OH- concentration pH paper = pH solution o pH < 7; acidic § Strong acids: 0-1 o pH = 7; neutral o pH > 7; basic § strong bases: 13-14 pH = – log [H+] pOH = – log [OH-] pH + pOH = 14 Ampholyte: can act as both an acid and a base pH as a function of concentration Actual pH vs Theoretical pH based from pH paper based from concentration of H+ in the solution prepared Relative Strengths of Acids and Bases Relative Acidity: HCl > CH3COOH > H2CO3 Relative Basicity: NaOH > NH3 Direction of the Reaction Direction of reaction is determined by the relative strengths of the acids and bases. The more acidic/basic the reactants, the more the reaction would proceed to react the reactants The reaction would proceed forward