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CHEM 1120: Final Exam

by: Tiana Roach

CHEM 1120: Final Exam 1120

Marketplace > East Carolina University > Chemistry > 1120 > CHEM 1120 Final Exam
Tiana Roach

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these notes cover everything that will be on the final exam.
Intro to Chem 1120
James Collins
Study Guide
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This 11 page Study Guide was uploaded by Tiana Roach on Wednesday April 27, 2016. The Study Guide belongs to 1120 at East Carolina University taught by James Collins in Winter 2016. Since its upload, it has received 38 views. For similar materials see Intro to Chem 1120 in Chemistry at East Carolina University.

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Date Created: 04/27/16
Chemistry Final Exam Study guide Reporting measurements  Measurements are reported to the limit of the measuring device plus a digit of  uncertainty. 1.20 cm  Graduated cylinder 1 digit of uncertainty Digital Balance 2 digits of uncertainty Thermometer 1 digit of uncertainty Buret 2 digits of uncertainty Significant Figures  Nonzero numbers 1­9  Leading zeros are not significant (0.003=1 sig fig)  Confined zeros are significant (109.505=6 sig figs)  Trailing zeros are significant (0.900=3 sig figs) In Calculations:  Add/subtracting: answer has same number of decimal places as the number with the  fewest decimal places.  Ex: 0.035+0.01=0.045, which is rounded to 0.05. The answer is  reported 2 decimal places because 0.01 has the fewest decimal places.  Mult/dividing: answer has same number of sig figs as number with the fewest amount of  sig figs. Ex: 5.356/1.63=3.285, which is rounded to 3.29. The answer is reported to 3 sig  figs because 1.63 has the fewest sig figs in the calculation. Avogadro’s number: 6.022*10 (use when converting between moles, atoms and molecules. Temperature Scales Celsius to Fahrenheit F=1.8*˚C+32 Fahrenheit to Celsius ˚C= F−32 1.8 Celsius to Kelvin K=˚C+273.15 Kelvin to Celsius ˚C=K­273.15 Heat Measurements  Heat is measured in Joules (J) or calories (cal).  1J=4.184 cal  To find heat: q=mCsΔT where m is the mass in grams, Cs is the specific heat and ΔT is the  change in temperature (T final­T initial). *this equation can be arranged to find specific heat, mass and change in temperature. Prefixes and conversion factors: Stoichiometry  Pay attention to amount of moles in stoichiometric equations.  When converting between units, start with given and work towards missing.  If units are cubed or squared, square or cube everything. 3 o How many cubic inches (in ) are in 10 cm? 1∈¿ 2.54 cm 3 ¿ 3. o ¿  ; The cm cancel out and the answer is expressed in in 3 (10cm )x¿ Periodic Table: Atoms, Naming Compounds and Assigning Charges  Atoms are the smallest particles  Subatomic particles includes protons and neutrons   The nucleus of an atom contains protons and neutrons.  Isotopes are atoms of the same elements with different  number of neutrons  Ionic Compound: metal + Nmetal (NaCl)  Covalent (Molecular) Nmetal + Nmetal  Charges of  ions are same as their group Number  Transition metals use roman numerals to represent charge  To find charges of atoms in polyatomic ions: 1) Crisscross Method. 2) Algebra; the sum  of all charges must equal zero.  Electronegativity: *fluorine is the most electronegative element, Polyatomic Ions, Molecular Compounds and Lewis Structures   Octet Rule: all atoms want to achieve electron configuration of a noble gas (8 valence electrons) Nuclear Reactions A B C  XE 1    E Y  2   Z 3 A  =  B  +  C  and  X  =  Y  +  Z The sum of mass number and atomic numbers must be the same on both sides of the  equation.  4  Alpha Decay: resembles Helium atom; 2He.  0  Beta Decay: resembles electron; neutron converts into proton and electron; ­1e or           0­1;   0   H +1e 0 ­1  Gamma Emission:  0                          1  Positron Emission: anti­electron; proton converts into electron and neutron; 1    n +   e 0 0 1  Electron capture : ­1  +   H1    n 10 2 I1 d2  Intensity:  I 2= 2   d1 Radionuclide stability  Radionuclides are stable when amount of neutrons are even and protons are even.  For lighter elements, a nuclide is most stable when the neutron to proton ratio, N/Z is  close to one. Half Life  Measure stability of radionuclides.  Longer half­life= more stable  Shorter half­life=less stable Fraction Remaining = 0.5n (where “n” = number of half­lives) Percentage Remaining = {0.5 } x 100%  Endothermic rxn: absorb heat from surroundings  Exothermic rxn: release heat to surroundings Endothermic S L                       Exothermic Intermolecular Forces of Attraction Dipole­dipole: permanent dipole, usually ionic compounds (transfer electrons); polar  (tetrahedral with all same peripheral atoms. London Dispersion: momentary dipoles; present in all molecules Hydrogen Bonding: usually molecular compounds (share electrons); exists between two  molecules of the same compound if have hydrogen atom covalently attached to a highly  3­  2­ ­ electronegative atom (N , O  and F). *Hydrocarbons are nonpolar and have similar electronegatives. *if molecule has hydrogen bonding, it has all three intermolecular forces.  Boiling points are influenced by intermolecular forces, molecular weight and molecular  shapes. Gas Laws Ideal Gas Law PV=nRT Boyles Law P1V 1P 2 2 Combined Gas Law P1V 1 P2V 2 = T 1 T2 Charles Law V 1 V 2 = T 1 T 2 2 1Pa= 1N/m 1Torr = 1 mmHg 1 atm = 760 mmHg R= 0.0821 L*atm/mol*K Gaseous Mixtures   Mole Fraction n A X A ¿nTotal  (moles of A÷ moles of gas)  % Composition= Mole Fraction X 100%  Daltons law: PtotalA+P BP C…. Solutions  Colloidal Dispersion: heterogeneous solution that does not settle out  Henry’s Law:  S1 = S 2 P1 P2  Saturated Solution: System at equilibrium; undissolved solute dissolves at same rate as  dissolved solute  Supersaturated Solution: unstable toward precipitating If Theoretical<Experimental; Supersaturated Expressing the Ratio of the Solute to Solution: Concentration Units: 1)  %(m/m)  =  [(mass of solute) / (mass of solution)] x 100 % 2)  %(v/v)  =  [(volume of solute) / (volume of solution)] x 100 % 3)  %(m/v)  =  [(mass of solute) / (volume of solution)] x 100 % 4)  ppm  =  [(mass of solute) / (mass of solution)] x  1,000,000 5)  ppb  =  [(mass of solute) / (mass of solution)] x  1,000,000,000 6)  Molarity  =  (moles of solute) /  Liters of solution  Dilution of solution1:  1    2 = 2M  x V  Osmosis: net movement of solvent from high concentration to low concentration; water  follows salt.  Hypertonic sol: increased solvent; water moves out; cell shrinks; crenation.  Hypotonic Sol: decreased solvent; water moves in; cell swells; hemolysis.  Osmotic Pressure: amount of pressure required to stop osmosis where M is  the molarity, R is the rate and T is the temperature.  Osmolarity: i  X Molarity, where I equals the sum of the coefficients Chemical Rxns, Energy, Rates and Equilibrium  Energy: the capacity to do work or transfer heat.  Kinetic Energy : energy in motion  Etotalkinetpotential m 2  Units of Energy: 1J=g 2  ; 1 cal= 4.184 J s  Bond energies can be used to estimrxn ΔH˚ =Σ[(reactants) ­  (products)]   Entropy: measure of disorder (ΔS). The more disorderly the more spontaneous. - In relation to phase changes:  Endothermic; orderly; nonspontaneous -ΔS S L G Exothermic; disorderly, sponΔSneous  Gibbs Free Energy: ΔG˚ = rxn ­ TΔrxn  rxn Spontaneity in relation to  Exothermic  Spontaneous                    Enthalpy, Entropy and  Exergonic (+ΔS) GFE                        (­ΔH) Endothermic Nonspontaneous(+ΔG) Endergonic                            (­ΔS)                             (+ΔH) Rates of Rxn  Rate =  ΔX  (rate of rxn depends on energy of activation). ΔTime  Lower energy= more stable  Exergonic rxn: energy released to surroundings; products lower in energy than reactants.   Endergonic rxn: energy absorbed by surroundings;; products higher in energy than  reactants.    Higher Energy of Activation= Progresses Slowly  Lower Energy of Activation= Progresses Rapidly  Catalysts speed up rxn by lowering energy of activation  Shorter the curve the faster the rxn Equilibrium Factors that affect Equilibrium 1. Molar Concentratioon a. Add/subtract amount of reactants and products. 2. Change in Pressure (only affects gases) a. Increase in pressure favors side with least amount of moles of gas 3. Change in Temperature - Equilibrium constant changes as temperature varies. a. First, Is the heat on product or reactant side? o +ΔH= endo, absorbed, reactant side o –ΔH= exo, released, product side b. Treat same as molar concentration (rxn will favor side w fewest moles) Bronsted­Lowry Perspective of Acids  Acids: increase hydrogen ion concentration (proton donor)  Bases: increase hydroxide ion concentration (proton acceptor) + ­  HCl  +  H O    H O   +   Cl    Conjugate acid­base pairs 2 3 HCl ­ (acid) (base)     (acid)     (base) /Cl pH, pOH, H 0 and OH Scales 3  NH   +   H O    NH +  +  OH­ 3 2 4 + (base)    (acid)      (acid)      (base)H O2 Interconverting between pH, +OH, H ­ and OH 3 - Six Equations for Interconvertin3 [H O ], [OH ], pH  and pOH: 1) K = [H O3] x [OH ] = 1.0 x 10 -14 w 2) pH = - log[H O 3 + + -pH 3) [H O ] = 10 3 - 4) pOH = - log[OH ] - -pOH 5) [OH ] = 10 6) 14 = pH + pOH Strong Acids and Bases 6 Strong Acids 8 Strong Bases The stronger the acid the weaker its conjugate base pair and vice versa. Salt Formation  Salt favors strong base or salt 1. NaOH + HCl= NaCl + H O 2 S.Base S.Acid Neutral Salt 2. NaOH + CH COO3= NaCH COO + H3O 2 S.Base W.Acid Basic Salt 3. NH +3 HCl= NH 4l W.Base S.Acid Acidic Salt Buffer Solutions  Resists large changes in pH by maintaining the H O Concentration 3  Consists of weak acids or weak bases. CANNOT be made from Strong acids or strong bases.  To find pH of Buffer Solution: A−¿ o If given K a pH=pKa+log( ¿ ¿ ¿ A−¿ o If given H O3 pH= -log(H O)3+ log( ¿ ¿ ¿


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