Chem 161 notes

CH1: Classes and Properties of Matter -Phases of Matter -Base Units & Measurements -Scientific Notation -Significant Figures -Density -Temperature(Conversion) -Dimensional Analysis Matter: Whatever Occupies Space and can be perceived by our senses. Pure Substances: Elements(Found on periodic table, cannot be broken down to a simpler substance; Sodium) & Compounds(Two or more Elements held together with a chemical bond; NACl) Mixture: Homogenous(Uniform Mixture, cannot be easily separated; broth) & Heterogenous(physically distinct parts that can be easily separated; soup) Physical Change: Change in the phase of matter but not the chemical identity, Liquid Water to Ice or Water Vapor/Steam, indicated by subscript, (s)(l)(g) - Solid Liquid Gas - in italics (aq) Aqueous Solution - Something Dissolved in Water, Solution - Solute / Solvent(h2o) NaCl(aq) NaCl(s) + H2O(l) -> NaCl(aq) Chemical Change: Change in the Chemical Identity that involves breaking or forming chemical bonds + chemical reactions Fe(s) + O2(g) -> FeO2(s) Base Units and Measurements, SI Base Units Mass: Kilograms(Kg), g,ibs,oz,tons Length: Meters(m), in,ft,yd,miles Temperature: Kelvin(K), F,C, Time: Seconds(S), Minutes, Hours, Years, blah blah blah Amount/Quantity of a substance: Mole(mol), 6.02*10^23 1 Megameter = 1mm = 1*10^6 m 1 Kilometer = 1 km = 1*10^3 m 1 decimeter = 1dm = 10^-1 1 centimeter = 1cm = 10^-2 1 millimeter = 1mm = 10^-3 1 micrometer = 1 MM = 10^-6 1 nanometer = 1nm = 10^-9 1 picometer = 1pm = 10^-12 Standard Notion -> Scientific Notation 73562 -> 7.3562 * 10^4 0.0000572 -> 5.72 * 10^-5 1000572 -> 1.000572 * 10^6 653->6.53*10^2 0.0157 -> 1.57 * 10^-2 0.0094->9.4*10^-3 Significant Figures - Set of rules to determine how numbers the answer should be reported to and why Insignificant Figures: Zeros before a Non Zero number and the decimal point. 0.034 - 2 0.0340 - 3 340 - 2 340.0 - 4 340. - 3 3401 - 4, write with 2 sf. 3400, write with 3 sf, use scientific notation, 3.40 * 10^3 3 parts to correct answer/calculation 1. The Math itself(duh) 2. The Units 3. Current Number of Significant Figures Addition/Subtraction - Determined by # with lowest decimal places, 27.4 + 6.92 = 34.32, 1dp, 34.3, only as much precision as the least precise value Multiplication/Division - Determined by # with the least amount of Significant Figures, 9.962g, 0.099860g/1ml, 4sf 5sf, have 4 sf, 9,962/0.099860 = 99.75966353, 99.76 mL, rounding rules apply 345-3 3450-3 642.0-4 13.000-5 1.4*10^5-2 1.0004*10^5-5 0.0302-3 300-1 300.-3 0.300-3 0.0300-3 Dimensional Analysis: Unit Conversion How many minutes in 675 seconds, 675/60 = ~11.3 minutes Always start with given value(s), Seconds to Minutes, 675 Seconds| |1 minute|60 seconds, How many years are in 6.51*10^4 sec? sec-minutes-hours-days-years, 6.51*10^4 * 1/60 * 1/60 * 1/24 * 1/365.25, 0.00206 years, 2.06 * 10^-3 years Temperature Conversions T(K) = T(C) + 273.15 T(C) = (5/9)(T(F)-32) 35.5 C to F and to K T(K) = 308.65 degrees T(F) = 81.7 58 F to C and to K T(C) = 14 T(K) = 287 Density: Mass and Volume D = m/v, D = g/ml, D = g/cm3, 1ml = 1cm^3 mass of 3.50 g and volume of 11.5ml, find density 3.50/11.5 = 0.304 g/ml mass of 1.5kg, dropped into beaker changed volume from 100 to 255 ml, find density 1.5 kg = 1500 grams 1500/155 =9.7g/ml Convert 70. mph to inches/second 70 * 5280 * 12 / 3600 = 1200 CH2: Subatomic Particles Protons(p+): Gives element its identity, element determined by # of protons in the nucleus, positive charge, contributes to atomic mass(1 p+ is 1 amu) Neutrons(n0): Neutral Charge, in nucleus, contributes to atomic mass(1 n0 = 1 amu) Electrons(e-): Negative Charge, orbit nucleus, does not contribute to atomic mass(1e- = 1/2000 amu), mass of electron is negligible. different number of neutrons in an atom is an isotope, different number of electrons in an atom is an ion. p table: C, 6, 12.0111; atomic symbol(name of element), atomic number(# of P+), amu nuclide symbol: supscr A subscri Z X; amu, atomic number, atomic symbol amu = sum of protons and neutrons red: gaseous, black: solid, blue: liquid, white with black outline: unstable and not found in nature N-14: p+(7) n0(7), N-15: p+(7) n0(8), isotopes of each other, differ in neutrons therefore have different atomic masses N-14 has abundance of 99.4%, what is the average atomic mass? Weighted Average: (0.994 * 14) + (0.006 * 15) = 14.006 = 14.01 amu Calculate average atomic mass of neon to 4 significant figures Ne-20 90.48, Ne-21 0.27, Ne-22 9.25 (20 * 0.9048) + (21 * 0.0027) + (22 * 0.0925) = 20.1877 = 20.19 amu AS, A#, AM, p+, e-, n0 Cl, 17, 35, 17, 17, 18 Ag, 47, 107, 47, 47, 60 Hg, 80, 200, 80, 80, 120 Ar, 18, 40, 18, 18, 22 Atoms are charge neutral(#+ = #-), p+ = e- Periodic Table ordered by atomic number Groups: Vertical Columns, Indicative of the number of valence electrons available to bond with Periods: Horizontal Rows, Group 1: Alkali Metals, Hydrogen the exception(does not belong anywhere) Group 2: Alkaline Earth Metals, Transition Metals Group 6: Chalcogens Group 7: Halogens Group 8: Noble Gasses Metalloid Staircase: Separates the two parts of the periodic table from metals to non-metals(right is non-metals) Ionic Bonds are between a metal and a non-metal Calcium is an Alkaline Earth Metal Calculations with Different Elements(atoms, molecules, moles) Amu = molar mass, mass of 1 mol of that substance(g/mol) Use Mm to convert between grams and moles Na = Avogadro's Number, 6.02*10^23 atoms/molecules per mol use Na to convert between atoms/molecules and moles. atoms/moles <Na> moles <Mm> grams find molar mass of water(h2o) sum of all amu of the substances 2(H) + O, 2.02 + 15.99, 18.03 Calculate molar mass of glucose(C6H12O6) 6 * (12.01) + 12 * (1.01) + 6 * (15.99) = 180.12 g/mol how many grams of Carbon in 16.5 mols of carbon? 16.5 * (12.01) = 198 grams how many atoms of Sodium in 10.5 mols of Na? 6.02*10^23 * 10.5 = 6.32*10^{24} atoms how many grams of glucose in 2.3 mols of glucose? 180.12 * 2.3 = 410 grams how many atoms of Hydrogen in 3.5 mols of Water? 6.02*10^23 * 3.5 * 2 = 4.2*10^{24} atoms how many grams are in 10.5 * 10^26 atoms of sulfur? 10.5*10^26 * 1/6.02*10^23 * 32.065 = 55927.3255814 grams compound: any kind of combination of elements molecules: covalent bond water: either, NaCl: Ionic, compound molarity = mol/l, unit of a concentration of solutions(sol^n) Volume L solute: less of solvent: more of(usually water(aq)) moles of solute / liters of solution, assume volume of solvent is the volume of the solution NaCl(s) + H2O(l) -> NaCl(aq) solute + solvent -> solution aq are homogenous mixture what is the concentration of NaCl(aq) when 55.0g of NaCl(s) is dissolved in 255ml of water? 55.0g * 1/(58.44) = 0.941 mol 255ml * 1 L/1000ml = 0.255L 0.941/0.255 = 3.69 M(olar) or mol/L Concentration when 55.0 grams of NaCl are dissolved in 755 mL of h2o? 0.941mol 0.755L 0.941/0.755 = 1.25 M increasing volume of solvent will decrease concentration of the solution how many grams of glucose are required to make a 250.0 mL of a 2.5M glucose solution? 2.5M = 0.625mol/0.2500L, 0.625 mol * 180.12g/1 mol = 110 grams 112.575g Accuracy: How close a measurement is to the true value Precision: How close measurements are to each other CH3: Electrons Protons(P+): Nucleus, Gives atom its identity. Neutrons(N0): Nucleus, contributes to Atomic Mass, Sum of P+ and N0 = AMU, isotopes of same element with different # of N0. Electrons(E-): Orbit Nucleus, do not contribute to atomic mass, Wave-Particle Duality, Ions form from different # of E- (Cations/Anions) W-P Duality: Have properties of both waves and particles, neither definition is complete. Electrons propagate through space as Waves, but they interact at a point like a particle. Waves: Wavelength(distance between 2 peaks/troughs) measured in nanometers(nm). Frequency (number of wavelengths that pass a fixed point in a unit of time) measured in 1/seconds or seconds^-1 or Hertz(Hz). Speed of Light(C) = 2.998 *10^8 m/s, C = Lambda * v (wavelength * frequency) What is the wavelength in nm of blue light with a frequency of 6.4*10^14 s^-1? 470 nm. What is the frequency of a wave having a wavelength of 681 nm? 4.40 * 10^14 s^-1. Energy: Joules(J) or Kilojoules(kJ), E = hV (Planck's Constant(6.626 * 10^-34 Js) * frequency) Energy of wave with frequency of 5.25 * 10^14? 3.48 * 10^-19 Joules Energy of wave with wavelength of 700 nm? 7.00 * 10^-7 m, 2.998/7.00*10^-7 = 4280000 or 4.28 * 10^6 s^-1, 6.626*10^-34 * 4.28 * 10^6 = 2.84*10^-19 Joules Energy Levels(n) n = 1, 2, 3, 4, etc etc. Can have energy being absorbed or emitted(added or removed) Electron transitions from a lower energy level to a higher energy level(added) +Delta-E NF = 3, Ni = 1 Electrons transition from a higher energy level to a lower energy level(removed) - Delta-E NF = 1, Ni = 3 Delta-E = -R_h * (1/nf^2 - 1/ni^2) Change in Energy(J/kJ), Rydburg constant(2.179 *10^-18 Joules), Final Energy Level, Initial Energy Level What is the energy and wavelength of light emitted when a photon transitions from n=6 to n=3? -2.179*10^-18 * (1/9 -1/36) = -1.816*10^-19 Joules, 1.816*10^-19 / 6.626 * 10^-34 = 2.74 * 10^14, 2.998*10^8/2.74*10^14 = 1.09 * 10^-6 M or 1090 nM Electrons as Particles Electrons and their interactions are responsible for chemical bonds and chemical reactions Quantum Mechanics: allows us to make a statistical statement(prediction) about the region in which we are most likely to find an electron. 4 Quantum Numbers for each electron is like its unique mailing address 516 High St. Bellingham, WA least to most specific n(principle quantum number)(size and energy level) Can be any positive integer starting from 1 l(angular momentum quantum number)(Shape of the orbital) Any integer from 0 to n-1 ml(magnetic quantum number)(Orbital orientation) Any integer +l to -l ms(spin quantum number)(Rotation of the Electron) +1/2 or -1/2 n=4, l = 0-3, ml = 3 2 1 0 -1 -2 -3, ms = +-1/2 possible sets n 1 3 5 2 l 1 2 3 1 ml 0 -1 -5 1 ms 1/2 -1/2 1/2 0 no yes no no set up on periodic table n going down from 1 to 7 set by how many electrons an orbital can hold s block room for 2 electrons p block room for 6 electrons d block room for 10 electrons f block room for 14 electrons S x P x x x D x x x x x F x x x x x x x Aufbau principle electrons will always fill the lowest energy level first Hunds rule: must have maximum number of unpaired electrons before you fill each orbital Draw orbital diagram of Oxygen(tell the location of the electrons) 1) How many electrons? 8 1s x 2s x 2p / / x Orbital diagram of potassium How many electrons? 19 1s x 2s x 2p x x x 3s x 3p x x x 4s / Electron Configuration: 1s2 2s2 2p6 3s2 3p6 4s1 Noble Gas Configuration: from the previous noble gas, [Ar]: 4s1 Orbital Diagram, electron configuration, and noble gas configuration of Bromine 35 electrons 1s x 2s x 2p x x x 3s x 3p x x x 4s x 3d x x x x x 4p / x x 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 [Ar]:4s2 3d10 4p5 Valence Electrons: Electrons in the outermost shell of an atom(highest energy level) that are available to bond. Only electrons available for chemical bonds. for example: Bromine only has 7 electrons available for bonding. Atoms are charge neutral(P+ = E-) ions are either positive(cation P+>E-, name + ion) or negative(anion P+<E-, suffix -ide), gaining/losing electrons to have a full outer shell. Atoms are most stable with a full outer shell. Octet except for hydrogen and helium. Bromine wants 36 electrons, needs one more. Bromide Ion still has 35 protons. Br- or Br-1. iso-electronic: same electrons noble gasses do not form ions, inert/non-reactive. Na has 11 p+ and e-, Na+ has 10 e- Aluminum has 13 p+ and e-, Al3+ has 10 e- Carbon, Silicon, and Germanium do not form ions. Sulfur, S2- Sulfide Calcium, Ca2+ Phosphorous, P3- Phosphide. Magnesium: Mg2+ Ionization Energy: The amount of Energy required to remove an electron from its outermost shell *number of e- in the outer shell *distance between the nucleus and the outer shell increases across a period and decreases down a group electron affinity: attraction of an element to add an e- atomic radius increases down a group and decreases across a period Cations: Smaller Anion: Larger Li 3p 3e, Li+ 3p 2e, AR decreases F 9p 9e, F- 9p 10e n=2 -> n=3 e- repels, AR increases Chapter 4: Chemical Bonds Electrons are responsible for Chemical Bonds Ionic Bonds: Compound formed between two ions, electrostatic attraction, Cation(metals) Anion(non-metals) Naming convention i.e. NaCl not ClNa, Sodium Chloride. Charge Neutral, balance the compound with subscripts. + = -. Covalent Bonds: Sharing Valence Electrons. Molecules/Compounds. Between two or more non-metals. Carbon Monoxide, Dioxide, Water. Lithium Sulfide: Li+ S2- Li2S Aluminum Bromide: Al3+ Br- AlBr3 Barium Iodide: BaI2 Ba2+ I- Magnesium Phosphide: Mg3P2 Mg2+ P3- Potassium Nitride: K3N K+ N3- Monatomic ions form from one element Polyatomic ions: ions of more than one element. 7 Pa Ions: Carbonate(CO3^2-), Ammonium(NH4^1+), Hydroxide(OH^1-), Nitrite(NO2^1-), Nitrate(NO3^1-), Phosphate(PO4^3-), Sulfate(SO4^2-) Sodium Carbonate: Na2CO3, Calcium Carbonate: CaCO3, Calcium Hydroxide: Ca(OH)2, Magnesium Nitrite: Mg(NO2)2, Ammonium Sulfate: (NH4)2SO4 Ionic Compounds: Formed by the Electrostatic Attraction of cations and anions(+ and -) Monatomic Ions: one atom, predict the charge of the ion based on how many electrons would be gained or lost to have a full outer shell.(nearest noble gas) Transition Metals: Ions of transitional metal are cations. Can form multiple ions. Charge of a transition metal ion must be given in the name with Roman Numerals. Ex: Fe(Iron), must be told charge, Fe+(I) Fe2+(II) Fe3+(III). Iron(I) Oxide, Fe2O Iron(II) Oxide, FeO Iron(III) Oxide, Fe2O3 CuCl2 Cu2+ Copper(II)Chloride CuCl4 Cu4+ Copper(IV) Chloride Fe(NO3)3 Iron(III) Nitrate Cu3P Copper(I) Phosphide Pb(SO4)2 Lead(IV) Sulfate chromium(VI) phosphate Cr(PO4)2 tin(II) nitrite Sn(NO2)2 Covalent Molecules: Sharing Valence Electrons in the Outer Shell. Between Two or more non metals. No charges to balance, multiple formulas, multiple ways to combine the same elements. CO, CO2. HO, H2O NO3 Nitrogen Trioxide Dinitrogen Trioxide N2O3 P2O5 Diphosphorous Pentaoxide Nitrogen Trihydride NO3 CH4 Carbon Tetrahydride Sulfur Dioxide SO2 Exceptions to Covalent naming rules: Diatomic Molecules. 7 of them, never found alone. Hydrogen: H2 Oxygen: O2 Nitrogen: N2 Fluorine: F2 Chlorine: Cl2 Bromine: Br2 Iodine: I2 Fe(NO3)3 Iron(III) Nitrate Ionic Bond Lead(IV) Sulfate Pb(SO4)2 Ionic N2O5 Dinitrogen Pentoxide Covalent Ammonium Phosphate (NH4)3PO4 Ionic N2 Nitrogen Covalent AgBr Silver(I)Bromide Ionic Trinitrogen Diphosphide N3P2 Covalent Lewis Structures: 2-D representation, skeletal structure of a molecule. Shows the bonds between atoms. Valence electrons, e- in the outershell, highest energy level, farthest from the nucleus. Covalent Bonds form to sharing Valence e- to give all atoms a full outer shell. Octet for 8 electrons bar H and He which have a duet. Draw the Lewis Structure of Nitrogen Trihydride NH3 Carbon Tetrafluoride CCl4 dihydrogen monoxide H2O Carbon Dioxide CO2 Nitrogen N2 Oxygen O2 Carbon Monoxide CO Lewis Structures of Polyatomic Ions. Carbonate CO3^2- 24 electrons Ammonium NH4^+ 5 + 4 - 1 = 8 electrons Hydroxide OH^- 1 + 6 + 1 = 8 electrons Nitrite NO2^- 5 + 12 + 1 = 18 electrons Nitrate NO3^- 5 + 18 + 1 = 24 electrons Phosphate PO4^3- 5 + 24 + 3 = 32 electrons Sulfate SO4^2- 6 + 32 + 2 = 40 electrons Resonance: Double Bonds are in all positions simultaneously. Need to draw all possible positions when doing lewis structures Formal Charge: used to assign a value to each atom in a LS to determine the most correct LS when more than one LS is possible. Most Correct is when the atoms have an FC of 0. FC = #val e- - [#unshared e- + 1/2 bonded e-] CO2: O=C=O, 6 - [4+2] = 0. 4 - [0+4] O-C==O, Electronegativity: The ability of an atom to draw an e- to itself within a bond. Not to be confused with Electron Affinity, which is drawing an e- to an atom to form an ion. Increases up and across the periodic table and does not include noble gasses. Fluorine has highest electronegativity. Difference in electronegativity(x) value determines bond type. Ionic: X>=2, Li(1.1)F(4.0) = 2.9 Polar Covalent: Unequal Sharing of Electrons, 0.4<X<2. O(3.5)-H(2.1) = 1.4, S- S+, partially positive/negative. <-+. Nonpolar Covalent: Equal Sharing of Electrons, X<=0.4. C(2.5)-H(2.1) = 0.4. Ionic: Electrostatic attraction between ions, metals and non-metals, + and -, Large Electronegativity Difference. Chapter 5: Models Lewis Structure(2d representation of how the atoms are bonded) vsepr model: valence shell electron pair repulsion. A 3d representation of how it looks. SN: Steric Number, regions of electrons around the central atom. EG: Electron Pair Geometry, The geometry of a molecule that does not differentiate between regions(e- pair or bond), determined only by SN. MG: Molecular Geometry, the geometry of a molecule that does differentiate between regions. Lone pair of electrons takes up more space than a bond. Nitrogen Trihydride vs Ammonium ion. Molecular Polarity Polar Covalent Bond: Separation of Charge, difference in electronegativity higher than .4 but lower than 2 Non-Polar Covalent Bond: No separation of Charges, difference in EN lower than .4. Lowercase Delta to indicate partial charge. Is this molecule polar? Must use M.G. Sum of vectors is zero, therefore non polar. CH4, Sn - 4, EG tetra, MG Tetra Assign polarity by drawing vectors. Vectors point towards the more electro negative atom. CO2. Vectors point towards the oxygen atoms. Non-polar. NH3, Polar. Ammonium, Non-polar. Water, Polar. Nitrite, Polar. Orbital Hybridization: The Averaging of Atomic Orbitals to generate new equal set of orbitals to form covalent bonds. CH4, sharing 8 e-. 2 of the e- orbit in an s shape, 6 of the e- orbit in a p shape. 4 hybridized orbitals. 4 sp^3 orbits. Dependent on SN. Single bond 0- = Sigma bond Pi bond CO2 has 2 of each N2 has 1 s and 2 pi. CH4 has 4 s. CH 6: Intermolecular Forces(IMF) Interactions between Molecules Attractive/Repulsive forces between Molecules Not Chemical Bonds(Intramolecular Forces) 3 Classes of IMF 1. London Dispersion Forces(Weakest) 2. Dipole-Dipole Interactions 3. Hydrogen Bonding(Strongest) LDF(van der waals interaction) All molecules have LDF(Polar or not), attractive force caused by an induced dipole(non-polar molecules only have LDF) No separation of charge, but opposite charges between molecules still attract. Therefore, temporary induced dipole. Dipole-Dipole Interactions(Polar Molecules Only) Separation of charge, one part partially positive and one part partially negative. Positive Dipole attracted to negative dipole. Hydrogen Bonding, Not a Bond, strong attraction between a Hydrogen that is covalently bonded to an O/N/F and a Second Molecule containing an O/N/F. Water. Oxygens are attracted to the hydrogens of other Water Molecules. Shown as Dashed Lines. NH3: All 3 Kinds. H2O: All 3 Kinds. NH4+: LDF, H-Bond, Not Dipole-Dipole interactions. NO2-: LDF, D-D, Not H-Bond. CO2: LDF only. How IMFs affect physical properties(boiling point & viscosity)(solubility & miscibility)(ion dipole effect/interactions) Boiling Point: Temperature(Measurement of Heat(Energy)) The Stronger/more IMFs a Molecule has, the Higher the Boiling Point. H2O: LDF, D-D, H-Bond, 212F H2S: LDF, D-D, Not H-Bond, 76F Viscosity: Resistance to Flow. The Stronger/More IMFs in a Molecule, the Higher the viscosity. Size and Shape of Molecules impacts IMFs, specifically LDF. C5H12 shapes: Pentane: CH3-CH2-CH2-CH2-CH3, BP of 36C. 2-Methybutane: CH3-CH2-CH(-CH3)-CH3, BP of 28 C. 2-2-Dimethypropane: C surrounding by 4 CH3, BP of 9 C. Impacted by Surface area, all are Non-Polar without H-bonds. Solubility: Solid and Liquid. If soluble, Solid dissolved in the Liquid. NaCl(s) + H2O(l) -> NaCl(aq). If not, does not dissolve. Like Dissolves Like i.e. Polar Liquid Dissolves Polar Solid and opposite is true. Miscibility: Liquid and Liquid. If miscible, Uniform Solution. If not, Immiscible, separated. Like mixes Like. Same as above. In chemistry, we call Ionic Compounds Salts. NaCl = Table Salt. LiCl, NaBr, NaI, etc etc. Ionic Compounds/Salts are Soluble in Polar substances. Ion-Dipole Interaction, Salts ionize in polar substances. H2O - NaCl. Say we have a lot of water. When NaCl or any salt in water, it breaks up into Na+ and Cl-. They find the negatively and positively charged parts of the water and go there. CH7: Stoichiometry Quantities in Chemical Reactions & Mol-Mol Ratio Reactants -> Products Chemical bonds of the reactants are breaking and the chemical bonds of the products are forming. K + Br react to form KBr. 2K+Br2->2KBr. balance chemical equations with coefficients subscripts to balance compounds 2 mol of K and 1 mol of Br2 for every 2 mol of KBr Carbon Disulfide reacts with Oxygen to form Carbon Dioxide and Sulfur Dioxide CS2 + 3O2 -> CO2 + 2SO2 1 mol of CS2 and 3 mol of O2 react to form 1 mol of CO2 and 2 mol of SO2 Magnesium and Hydrogen Chloride to Form Magnesium Chloride and Hydrogen Mg + 2HCl -> MgCl2 + H2 1 mol of Magnesium and 2 mol of Hydrogen Chloride react to form 1 mol of Magnesium Chloride and 1 mol of Hydrogen Iron reacts with Hydrogen Sulfate to form Iron(III) Sulfate and Hydrogen 2Fe + 3H2SO4 -> Fe2(SO4)3 + 3H2 2 mol of iron and 3 mol of Hydrogen Sulfate react to form 1 mol of Iron(III) Sulfate and 3 Mol of Hydrogen CH4 + 2O2 -> CO2 + 2H2O SiO2 + 4HF -> SiF4 + 2H2O Mg(OH)2 + 2HCl -> MgCl2 + 2H2O Limiting Reactant: Reactant that runs out first. Excess Reactant: Reactant Left over that didn't react. Theoretical Yield: Amount of Product produced w/ the L.R. Actual Yield: Amount of Product actually made. Percent Yield: Actual/Theoretical * 100. 1a 3b 2c -> 1d have 7 a 10 b 12 c, Excess reactant: 4a 1b 6c Limiting Reactant: b, the ham Theoretical Yield: 3.33 3/3.33 * 100 = 90% 3g of Mg reacts with 2.2g of O to produce MgO, what is the limiting reactant, what is the TY, what is the PY if 4.5 of MgO is produced 2Mg + O2 -> 2MgO 3g of Mg, need MgO, from g to mol to mol to g, using molar mass and mol-mol ratio. 3/24.31g * 40.30g/mol = 4.97g of MgO 2.20/31.98g * 2 * 40.30g/mol = 5.54g MgO Magnesium is the Limiting Reactant. Can only produce 4.97 MgO 4.5/4.97 * 100 = 91% .23g of unreacted O2. 5g of H reacts with 17g of N to form 15.5 of NH3 is produced. 3H2 +N2 = 2NH3 5/2.02 * 2/3 * (14.007 + 3*1.0080) = 28.1 g of NH3 17/28.014 * 2 * (14.007 + 3*1.0080) = 20.7 g of NH3 Nitrogen is the limiting reactant TY: 20.7 PY: 15.5/20.7 * 100 = 74.9% 20.7*1/17.04 * 3/2 * 2.016 = 3.68g of H2. 5-3.68=1.32g of excess H2. Ca(OH)2 neutralizes HCl to produce CaCl2 and H2O Ca(OH)2 + 2HCl = CaCl2 + 2H2O 13.2g reacts with 10ml of 6M HCl Solution 13.2/(40.078+32.98+2.02) * 1/1 * (40.078+(2*35.453)) = 20.g 10/1000 * 6.00 * 1/2 * (40.078+(2*35.453)) = 3.3g HCl is the limiting reactant, Ca(OH)2 is the excess. 3.3*1/(40.078+(2*35.453)) * 74.08 = 2.2g of excess Calcium Hydroxide. Empirical formula: formula of a substance with the smallest integer subscripts molecular formula: formula of how the substance actually exists. Some whole number multiple of the empirical formula. C6H12O6, Molecular Formula CH2O, Empirical Formula Whole number integer: Mf/Ef in terms of molar mass. Solve for Mf from Ef. Benzene has e. formula of CH, molecular weight(Mm of Mf) is 78.1 g/mol, what is Mf of Benzene 12.01+1.01 = 13.02 78.1/13.02 = ~6 C6H6 percent composition: mass percentage of each element in a compound Hexamethylene: 62.1% Carbon, 13.8% Hydrogen, 24.1% Nitrogen. What is the Ef of Hexamethylene? Assume 100g, convert % to g convert g to mol divide by smallest # of mol to get subscripts 62.1g of C, 13.8g of H, 24.1g of N 5.171, 13.69, 1.720 divide all 3 by 1.720 3,8,1 C3H8N1 Sodium Pyrophosphate 34.5% Na, 23.37% P, 42.1% O 34.5 g of Na, 23.3% g of P, 42.1% g of O 34.5/22.98, 23.3/30.96, 42.1/15.999