Chem 131, first exam
Chem 131, first exam CHEM131
Popular in Chemistry I - Fundamentals of General Chemistry
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This 9 page Study Guide was uploaded by Christina Notetaker on Friday September 23, 2016. The Study Guide belongs to CHEM131 at University of Maryland - College Park taught by John Ondov in Fall 2016. Since its upload, it has received 62 views. For similar materials see Chemistry I - Fundamentals of General Chemistry in Chemistry at University of Maryland - College Park.
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Date Created: 09/23/16
Chem 131 (chapters 1-4) First Exam Study Guide Intro: Here is a study guide that covers what we learned during Prof. Ondov’s lecture and topics from the syllabus that I gleaned from the textbook covering chapters 1-4. I’m basing this off of the study guide outline he posted on Canvas. I hope this helps. Chapter 1 ➢ Equations for forming positive or negative ions eg. 1 Lithium has the atomic number 3 for 3 protons and 3 electrons in its neutral stat n order for an + − ion to carry a positive charge, it must lose an electron: Li → Li + e (in other words, a positive ion or cation is formed from a neutral atom as taking away an electron leaves a more positive charge on the ion) eg. 2 Fluorine has the atomic number 9 for 9 protons and 9 electrons in its neutral state. In order for − − an ion to carry a negative charge, it must gain an electron: F + e → F (similarly, a negative ion or anion is formed from a neutral atom as gaining an electron leaves a more negative charge on the ion) List of Common C ations 3+ 4+ aluminum Al lead (IV) Pb ammonium NH 4+ lithium Li + barium Ba 2+ magnesium Mg 2+ 2+ 2+ calcium Ca manganese (II) Mn caesium Cs + mercury (I) Hg 2+ 2 chromium (II) Cr 2+ mercury (II) Hg 2+ 3+ 2+ chromium (III) Cr nickel (II) Ni + + copper (I) Cu potassium K copper (II) Cu 2+ silver Ag + + + hydrogen H sodium Na iron (I) Fe 2+ tin (II) Sn 2+ iron (II) Fe 3+ tin (IV) Sn 4+ 2+ 2+ lead (II) Pb zinc Zn List of Common A nions − 3− acetate CH3COO nitride N − − bromide Br nitrite NO 2 carbonate CO 2− oxalate C O 2− 3 2 4 hydrogen carbonate HCO 3 oxide O 2− − − chlorate ClO 3 permanganate MnO 4 − 3− perchlorate ClO 4 phosphide P chloride Cl − phosphate PO 3− 4 chlorite ClO − hydrogen phosphate HPO 2− 2 4 hypochlorite ClO − dihydrogen phosphate H P2 4− 2− − chromate CrO 4 sulphate SO 4 2− − dichromate Cr 2 7 hydrogen sulphate HSO 4 cyanide CN − sulphide S 2− fluoride F − hydrogen sulphide HS − − 2− hydride H sulphite SO 3 − − hydroxide OH hydrogen sulphite HSO 3 iodide I − thiocyanate SCN − nitrate NO 3− thiosulphate S 2 32− ➢ Calculate atomic mass from isotope masses and abundances eg. 1 Silicon consists of 3 isotopes, Si-28, Si-29 and Si-30, whose atomic masses are 27.9769, 28.9765 and 29.9738 respectively. The most abundant isotope is Si-28 which accounts for 92.23% of naturally occurring silicon. Given that the observed atomic mass of silicon is 28.0855 calculate the percentages of Si-29 and Si-30 in nature. a) (27.9769 amu)(0.9223) + (28.9765 amu)(x) + (29.9738 amu)(y) = 28.0855 amu x + y + 0.9223 = 1.0000 y = 1.0000 − 0.9223 − x y = 0.0777 − x (solve for y, then substitute) b) (27.9769 amu)(0.9223) + (28.9765 amu)(x) + (29.9738 amu)(0.0777 − x) = 28.0855 amu c) x = 0.0466 *00% = 4.67% Si − 29 y = 1.0000 − 0.9223 − 0.0466 y = 0.03101 100% = 3.10% Si − 30 * eg. 2 Rhenium has two naturally occurring isotopes: Re-185 (37.40%) and Re-187 (62.60%). The sum of the masses of the two isotopes is 371.9087. Calculate the atomic weights of Re-185 and Re-187. a) average atomic mass of Rhenium- 186.207 amu b) (x)(0.3740) + (y)(0.6260) = 186.207 amu y = 371.9087 − x (substitute) c) (x)(0.3740) + (371.9087 − x)(0.6260) = 186.207 amu d) x = 184.95 amu Re − 185 e) (371.9087 amu Re) − (184.95 amu Re − 185) = y y = 186.96 amu Re − 187 ➢ Write Isotope symbols. Determine the numbers of electrons, protons, and neutrons in an isotope’s atom. Isotope symbol: mass number = number of protons + number of protons atomic number = number of protons = number of electrons ➢ Sketch a mass spectrum of a naturally occurring sample of an element given the abundances and masses of its various isotopes. >2 peaks- 2 isotopes >most common isotope is at 194 amu due to high relative abundance ➢ Solve problems by applying the Law of Definite Proportions. Solve problems involving the Law of Multiple Proportions Law of definite proportions- a sample has the same proportions as the elements that compose it regardless of how it is created eg. 1 + 2− 1 * O 2 2 H +*1 O * (any chemical compound will contain a fixed ratio of elements by mass) eg. 2 + 2− 5 * O 2 10 H +*5 O * for every water molecule, there are 2 H atoms and 1 O atom Law of multiple proportions- two elements that are combined to form two different compounds produce two masses; compounds are expressed in a small ratio of whole numbers eg. 1 C + O 2 CO, where 1g C and 1.333g O produce 2.333g CO C + O 2 CO , w2ere 1g C and 2.333g O produce 3.333g CO 2 Chapter 2 ➢ Use equations for volume, density, and kinetic energy to solve for one variable given the other variables. V = base x height density = mass/volume KE = 1/2mv 2 ➢ Memorize the base metric units and the metric prefixes. Convert from one to another. SI base units (- “systeme internationale” or International System of Units) Quantity Unit Symbol length meter m mass kilogram kg time second s electric current ampere A amount of substance* mole mol temperature Kelvin K luminous intensity candela cd *chemical substance that contains ions, atoms, molecules, protons, etc. SI prefixes T G M k h da | d c m μ n p Tera giga mega kilo hecto deca | deci centi mili micro nano pico 1012 109 106 103 10 10 | 101 102 10-3 10-6 10-9 10-12 The great monkey king has died | drinking chocolate milk, μoving nine palaces Other conversions: 3 1 L = 1000cm 1 L = 0.001 mL 3 1 mL = 1 cm ➢ Be able to obtain derived units for volume and area Common Derived SI Units 2 area m volume m (cubic meter) density kg/m 3 energy J or kg m /s (Joules) * frequency Hz or s−1 (Hertz) 2 pressure Pa or n/m (Pascal) electric charge C (Coulomb) ➢ Convert between number of moles and number of atoms; and mass (in grams or kg), number of moles, and numbers of atoms using Avogadro’s Number and atomic mass of an element. -23 -1 Avogadro’s number- 6.022 x 10 mol moles to atoms/molecules: (number of moles) (6.022 x 10 −23) = number of molecules * moles to mass to molecules/atoms: mass molar mass = number of moles unit form: (g/mol) = (grams) (moles) Chapter 3 ➢ Convert from frequency to wavelength; wavelength to frequency for electromagnetic radiation. wavelength to frequency: c= f (c-speed of light, 3 x 10 m/s) λ ➢ Sketch the apparatus used to demonstrate the photoelectric effect. ➢ Understand the variables in the photoelectron energy equation (KE = hv - ϕ) and how to use it. Understand the concept of threshold frequency. KE = hv − ϕ KE- kinetic energy, the difference between energy of photon and binding energy of electron E final hv v- velocity, speed or frequency of an electron h- Planck’s constant, 6.626 x 10-3J*s ϕ = hv = E initial ϕ- binding energy, energy binding electron to metal, symbol Փ ➢ Make calculations using the following equations: Bohr’s equation for the energy of the electron in the H atom, de Broglie’s equation for the wavelength of a particle, Heisenberg’s uncertainty equation. Bohr’s equation: 1 1 ΔE atom = − R ( h 2 − 2 ) (R- Rydberg constant, 2.18 x 10 -18J) n f n i de Broglie’s equation: h -34 22 λ = mv (h- Planck’s constant, 6.626 x 10 Js or kg m /s s) Heisenberg’s uncertainty equation: Δx × mΔv ≥ h 4π ➢ Sketch orbital shapes (s, p, and d); write their quantum numbers (n, l, ml) Atomic Orbital Shapes: s- sphere p- dumbbell d- clover leaf ➢ Sketch radial probability distributions for ns, np, and d orbitals of the 3rd period. probability density- probability of finding an electron at a point in space radial distribution function- total probability of finding an electron within a spherical shell Chapter 4 ➢ Write electron configurations for s & p block elements and for transition elements of periods 4 and 5, and their ions. eg. 1 K (atomic number- 19): 1s 2s 2p 3s 3p 4s ⇒ [Ar] 4s 1 eg. 2 As (atomic number-33): 2 2 6 2 6 2 10 3 2 10 3 1s 2s 2p 3s 3p 4s 3d 4p ⇒ [Ar] 4s 3d 4p eg. 3 Ru (atomic number- 44): 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d ⇒ [Kr] 5s 4d6 2 6 eg. 4 Hg (atomic number- 78): 10 10 14 8 14 8 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d ⇒ [Xe] 6s 4f 0 5d ➢ Know the charges on those ions that have predictable charges. (see common cations and anions under chapter 1) ➢ Determine the number of valence electrons in atoms in the s and p blocks and the first two transition metal series. eg. 1 Mg: 2 valence electrons S: 6 valence electrons Sc: 2 valence electrons, due to the 4s2 orbital being greater than the 3d1 Ti: 2 valence electrons, ^^ ➢ Place in order the atomic radii of (a series of 3) s & p block atoms and their ions. effective nuclear charge- net charge an electron experiences atomic radius increases down a group/column ( each energy level is further from the nucleus) and decreases across a period/row (there is a higher effective nuclear charge since there are more protons, creating a greater attraction force pulling the electrons toward the nucleus) ➢ Write the defining equations for 1 t, 2nd,3rd, . . . ionization energies eg. 1 + first ionization energy of Al: Al (g) → A (g) + e − + 2+ − second ionization energy of Al: Al (g) → Al (g) + e third ionization energy of Al: Al (g) → Al (g) + e − 3+ 4+ fourth ionization energy of Al: Al (g) → Al (g) + e − ➢ Write the defining equation for the electron affinity of an atom. electron affinity- energy emitted when a neutral atom gains an electron in the gas phase electron affinity decrease down a group/column (the added electron is placed in a larger orbital, which is farther from the nucleus and increases the electron-electron repulsion force) and increases across a period/row (the added electrons are on the same orbital shell and are closer to the nucleus) ➢ Write a full set of proper quantum numbers for an electron in an orbital specified. eg. 1 Oxygen, 2p : n=2, l=1, ml=0, ms= +1/2
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