Chem 1 Lab - Summary Lab For Chem 1211 PDF

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All notes from all labs between chapters 1-10 with Professor Edward Witten....

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Lab 1 1. Accuracy vs Precision 1. Precision: refers to the reproducibility of our measurements – how close are our measurements to each other 1. Take multiple repetitions and take the average of them all 2. You can be precise and not accurate, but not visa versa 2. Accuracy: refers to how close our measured value is to the actual value or the “accepted” value 1. An “accepted” value is a value that has been determined experimentally (so, it has some uncertainty in it) and is generally accepted by the scientific community 2. Significant figures 1. All non zeros are significant 2. Zeros between non zeros are significant 3. Zeros after a decimal point to the right are always significant 4. Zeros to the left of the first non zero number are not significant 5. In addition and subtraction: same number of decimal places as the fewest one 6. In multiplication and division: same number of significant figures overall as the fewest one 3. Density: d=m/v 1. Density is a physical property of matter, an intensive property 2. If we place a solid substance in a liquid of higher density, the solid will float on the surface of the liquid. If the density of the liquid is less than that of the solid, the solid will sink. If the density of the liquid is the same as the density of the solid, the solid will “just float.” 3. Water displacement technique 1. Can determine density by taking the mass of object, finding an initial volume of water, dropping object in and recording new volume, then subtract initial volume from final volume to get the volume to use in the equation 4. Reading graduated glassware 1. Always read the bottom of the meniscus at eye level 2. Estimate the last significant figure 5. Design an experiment in order to separate compound A from compound B. List all the steps. 1. Filtration 2. Decantation 1. Separating a liquid from a solid (centrifugation) 3. Extraction 1. Dissolve a component of a mixture into a solvent 2. Then evaporate the solvent to get your component 4. Separation funnel 1. Immiscible liquids Lab 2 1. Molar mass and avagadro’s number conversions (6.022 X 10^23 particles) 1. Represents the number of atoms in 12 grams of carbon-12 2. Unit conversions 3. Molar mass 1. The molar mass of a substance is the sum of the molar masses of each element

Lab 3 1. Atomic emission spectra: 1. Consist of bright lines at discrete wavelengths on a dark background 2. Produced when free gaseous atoms are heated 3. Hydrogen Gas 1. Four bands; red (highest wavelength), green, blue, violet (lowest wavelength, most energy) 4. Colors: Red (3-2), green (4-2), blue (5-2), violet (6-2) 5. Atoms heated in a flame emit light consisting of bright lines of discrete λ (orange light) 2. Atomic absorption spectra: 1. Consist of characteristic series of dark lines in continuous spectra 2. Produced when free gaseous atoms are illuminated by external source of radiation 3. Electron movement causes light 1. When electrons gain energy they get excited and jump up an orbital 2. When electrons go down an orbital energy is released and color is seen 4. E = hc/wavelength (make sure to convert wavelength into meters) (10^-9 m) 5. Flame Test: 1. Electron absorbs energy from the flame goes to a higher energy state 2. Electron goes back down to lower energy state and releases the energy it absorbed as light 6. Sodium can be recognized by a bright orange flame color Lab 4 1. Know how to draw lewis structures 1. Add up valence electrons 2. Put atom with highest bonding capacity in the center 3. Obtain octets on all atoms (might need double or triple bonds) 4. Place lone pairs on center atom if any 2. VSEPRTheory: based on minimizing repulsion between pairs 1. Know electron and molecular geometries and bond angles 2. Steric number (SN) = number of atoms bonded to central atom + number of lone pairs on central atom 3. Non polar Covalent bond: two non metals share equally (symmetrical around central atom) 4. Polar Covalent bond: two non metals share unequally (asymmetrical around central atom) 5. Shape of a molecule can effect solubility, physical state, and chemical activity 6. Ionic (Nonmetal and a Metal): 1. Metals lose valence electrons to achieve noble gas electron configuration 2. Nonmetals gain electrons to achieve noble gas electron configuration 7. Allotropes: 1. Different molecular forms of an element 2. Ex. Oxygen: O2/O3 (ozone) 8. Formal charge (FC): 1. Measure of the number of electrons formally assigned to an atom in a molecule 2. FC = (# valence electrons) – [ # nonbonding electrons + 1/2(# bonding electrons)] 9. Exceptions to the Octet Rule 1. Molecules with atoms having fewer than eight electrons = electron deficient 2. Molecules having odd number of electrons = free radicals (very reactive) 3. Molecules with atoms having more than an octet = expanded valence shell 10. Electronegativity (EN) trends (greatest in top right corner)

Lab 5 1. Chromatography 1. Analytical technique used to separate mixtures into their component parts in order to analyze, identify, purify, and/or quantify the mixture or components 2. Chromatography separates components within a mixture by using the differential affinities of the components for a mobile medium and for a stationary adsorbing medium through which they pass 3. Differential affinities are dictated by intermolecular interactions between the components of the sample and the stationary and mobile phases 2. Mobile phase: gas or liquid that carries the components of the sample 3. Stationary phase: the part of the apparatus that does not move with the sample 4. Capillary Action: the movement of liquid within the spaces of a porous material due to the forces of adhesion, cohesion, and surface tension. The liquid is able to move up the filter paper because its attraction to itself is stronger than the force of gravity 5. As the water moves through the filter paper, the components of the mixture that are soluble are dissolved and move with the solvent. The components that are attracted strongly by the water and weakly by the paper will move rapidly, while the components that are attracted more strongly to the paper will move very little and slowly. 6. Don’t let the sides of the paper touch or else the dyes will move bidirectionally 7. Finding Rf values 1. Rf = distance traveled by spot/distance traveled by solvent Lab 6 1. Calculating percent out of hydrate (%x= mass x/mass compound) 2. Chemical formula of a hydrate (MxAz x nH2O) 1. M is a metal cation 2. A is an anion 3. n is number of water molecules 3. Hydroscopic: compounds that absorb water from the atmosphere 4. Anhydrous: compound that is left after the water is taken out 5. Dehydration and Rehydration 6. %x = Mass x/ Mass Total 7. Find formula of the hydrate (empirical formula) Lab 7 1. More reactive metals will replace less reactive metals 2. Calculate percent theoretical yield, actual yield, and percent yield 1. Percent yield = actual yield/theoretical yield x 100 3. Calculate limiting reactant (one with less) Lab 8 4. Titration: determine the concentration of an unknown reagent by reacting it with another reagent of known concentration. 5. Equivalence point: when enough standard solution has been added to completely react with the unknown substance 6. End point: when indicator that was added changes color 1. Phenolphthalein turns pink in basic solutions

7. Fermentation: food preservation method by using lactic acid bacteria 8. Calculate percent lactic acid 9. Determine concentration of a solution (Macid·Vacid = Mbase·Vbase) 10. You need to know the concentration of one of the two solutions in order to use titration methods to find the concentration of the other or else all you will get is a ratio 1. Ex. You want to determine the concentration of a solution of HCl. You titrate 5.0 mL of HCl with 0.9581 M NaOH. You used 4.96 mL of NaOH to titrate the HCl. What is the concentration of your HCl solution? 11. Components: 1. An indicator: phenothaline which will change to a pink color in a basic solution 2. Solution of known concentration called the titrant 3. Solution of unknown concentration 4. Graduated cylinder, conical flask, burette 12. Mole Ratio: Moles Acid/Moles Base 1. Find Ratio and if not 1:1 then cross multiply (Mol A/Mol B = ?/?) 2. Turn it into MaVa=MbVb 3. Rearrange to find variable you are looking for Lab 9 1. When an ionic compound dissolves in water, it dissociates into its constituent ions 2. When a solute dissolves in a solvent 3. Precipitation reactions: 1. Occur when two solutions of ionic species are mixed and a solid product (precipitate) is formed. Precipitation reactions are examples of double replacement reactions or metathesis reactions 2. KA+MB→ KB+MA 4. Solubility Rules: 1. All common compounds of the Group 1 metals (i.e. Na+, K+, etc.) and of the ammonium ion, NH4+, are soluble in water 2. All common compounds of halides (Cl-, Br-, I-) are soluble in water except the halide salts of Ag+, Hg22+, and Pb2+ 3. All common nitrates, NO3-, are soluble in water 4. All common sulfates, SO42-, are soluble in water except PbSO4, Hg2SO4, and BaSO4. Slightly soluble sulfates include CaSO4, Ag2SO4, and SrSO4 5. The important soluble metal hydroxides are NaOH and KOH. Most other metal hydroxides are insoluble. Slightly soluble metal hydroxides include Ca(OH)2, Ba(OH)2, and Sr(OH)2 6. All common carbonates, CO32-, phosphates, PO43-, and sulfides, S2- are insoluble in water except those salts of the Group 1 metals and NH4+. (The sulfide salts of Group 2 metals are also soluble) Lab 10 1. Heat is defined as energy transferred from one system to another because of a difference in temperature 2. Heat always flows from the hotter to the colder system 3. -qhot = qcold 4. All chemical reactions either generate (exothermic, -) or absorb (endothermic, +) heat

5. When reactions are conducted at constant pressure (for example, in the lab under atmospheric pressure) heat of reaction is called enthalpy of reaction and is given the symbol ∆H 6. HESS’s law 7. Calculation of ∆Hs from temperature change data...