Gen Chem 2 Final Review PDF

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Dr.Gavva entire Chem 2 course overview for Final Exam...

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General Chemistry 1 - Final Review, Fall 2014 Chapter One - sig figs, unit conversions, metric conversions  (6.35+3.6)/99.5 = 0.10 – Least Sig figs (for multiplication/division)  5.3 + 5.44444 = 10.7 – Least decimal places (for addition/subtraction) Chapter Two - parts of the atom  Proton, neutron, electron. Protons MUST EQUAL the electrons in the ground state of an atom. - atomic number, mass number, isotopes, atomic symbols - the mole and molar mass - A = mass number - Z = atomic number - Nuclide is the element - Quantum Numbers: o n, l, ml, ms o n = energy level, ranges from 1 to … o l = orbital shape, n-1. (remember s, p, d, f) and ranges from 0 to n-1 o ml = angular quantum number referring to orbital orientation.  Ranges from –l to l (L not 1) o ms = magnetic spin refers to electron spin in an orbital.  Can only be +1/2 or -1/2 for an individual electron - There 2 electrons in each orbital o Which is why 1s2 means there are 2 electrons in the 1s orbital and is represented by a single box - Remember there are two ways for electron configuration o Noble Gas Configuration: Mg: [Ne]3s2 o Electron Configuration: Mg: 1s22s22p63s2 - Diamagnetic v. Paramagnetic: o Diamagnetic have no partial spin, so either the orbital is full or empty. o Paramagnetic have some partial spin, so there is one electron in the orbital. (that 1s orbital picture above would be sort of an example) - All Trends across Periodic Table o Zeff = effective nuclear charge that holds together all the electrons  This is why Cations are smaller than Anions since the number of protons outweigh electrons and have a stronger pull against them which is why the electron cloud/ electrons come closer making the atomic radius smaller.  If you have isoelectronic elements; the higher charged cations will be the smallest atoms and the higher charged anions will

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be the largest. (isoelectronic means they have the same number of electrons) o IE = Ionization Energy = Energy needed to remove an electron o EA = Electron Affinity = The “Affinity” or attraction to gain an electron. Elements like N, O, F have 3-, 2-, 1- charges have increasing EA AND IE (Look at the trends at the bottom of the page). o Remember: All atoms want to be in the most stable position or reach an “octet” (8 Electrons) o Electronegativity is relative to each element, with F have the largest (4.0) and Fr have the smallest (0.7) Rutherford – Gold foil – Discovered Protons or “Alpha Particles” Thompson – Cathode Tube Ray – Discovered Electrons Chadwick – Discovered the Neutron Millikan’s Oil Drop Experiment – determined charge of electron

Chapter Three - energy of a photon  (wavelength = (ch/∆E), ∆E = hv, ∆E = ((R/n²f)-(R/n²i))); R = -2.178x10-18  C = λν, where c = 3.00x108 m/s - quantization of energy and Bohr equation

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Chapter Four - periodic table  1A = Alkaline  2A = Alkali Earth  6A = Chalcogens  7A = Halogens  8A = Noble Gasses - periodic trends  See the attached material at the end of this packet.

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Valence Electrons determined by group number. (i.e. F is in group 7A so it has 7 valence electrons.

Chapter Five - Nomenclature o Divided into two groups  Metals and Nonmetals are IONS with the metals as the cation and nonmetals as the anion. Metal is normal but the Nonmetal ends with an –ide (e.g. sodium chloride)  Nonmetals with Nonmetals follow similarly but the prefixes for indicating how many there are apply. (i.e. P2O5 is diphosphorous pentoxide, BUT if it’s only one for the first element you do NOT say mono-. i.e. CO2 is carbon dioxide not monocarbon dioxide.) o Acids  Let’s just analyze one case and then most of the others become apparent:  The Chlorine Acids: o HClO – (hypo)chlorous acid o HClO2 – chlorous acid o HClO3 – chloric acid o HClO4 – perchloric  The same structure applies to other halogen families (i.e. Br and I)  HS – hydrosulfuric acid  HCl – hydrochloric acid o (note the hydro in these cases)  Transition metals oxidation states: o Fe has +2 and +3 o Pb has +2 and +4 o Zn has +2 o Cu has +1 and +2 o Mn has +2 and +4 o Ni has +2 o Cr has +2 and +3 o Ti has +4 o Co has +2 and +3 - ionic and covalent bonds - An ionic bond is solely based on charge and attraction. No electrons are actually shared. Covalent bonds actually share their electrons. Ionic bonds are always between a metal and a non-metal, whereas molecular (covalent) bonds are between two non-metals. - molar mass

Chapter Six - Lewis structures

- electronegativity and polarity Remember that dots are one and sticks are two for formal charge. Any compound that is made up of non-like elements (anything except the BrINClHOF elements) will have dipole-dipole movement. Anything that has an H bound to an O, F, or N will have hydrogen bonding. Note that CH2O does NOT have hydrogen bonding because the hydrogens will be bound to the carbon, not the oxygen. Chapter Seven - VSEPR theory and molecular geometry  VSPER is the prediction of molecular shape based on the lewis diagram. Do NOT confuse molecular for geometric. - molecular geometry and polarity - intermolecular forces - Formal charge = Valence electrons – ½ bonded electrons – lone pair electrons

Chapter Eight - balancing chemical equations  Remember to balance lone atoms last (ie. HCl + O2 -> ...) - stoichiometry and limiting reagents - Use Dimensional Analysis

o Don’t think about whether you should divide or multiply to get the units you want; let the unit conversions do the work for you by cancelling out the units you want in order for you to get the units you need. I know this sounds weird but it’s how it should be. Chapter Nine - balance simple redox  OIL RIG or LEO says GER (Oxidation Is Loss - Reduction Is Gain, or Loss Equals Oxidation - Gain Equals Reduction) - Oxidation states you should know: o O has a 2o Halogens have 1o Peroxide (O22-) has a 1-

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predict precipitates

Chart:

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Always NAG people who SAG except for when they PMS or CaStroBar

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N - Nitrates A - Acetates G - Group 1

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S - Sulfates (†*) A - Ammonia G - Group 7 (†)

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† Except for when they (have)

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P - Lead (Pb) M - Mercury (Hg) S - Silver (Ag)  or

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Ca (Calcium) Stro (Strontium) Bar (Barium).

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NAG SAG states things that are always soluble except in the presence of the exceptions. The exceptions always form precipitates.

Strong bases are anything that have a hydroxide (-OH) group attached. Strong acids are H2SO4, HBr, HI, HCl, HClO3, HClO4, HNO3. All strong bases and acids are strong electrolytes. Something only has to dissociate with itself to become an electrolyte. Molecules like glucose are nonelectrolytes; they do not dissociate. Chapter Ten - heat, heat, and more heat ∆U = q + w. Be mindful of the terminology. If work is done ON a system, it is POSITIVE. Anything synonymous with this is POSITIVE. If else, it is NEGATIVE. - calorimetry ∆q(system) = -∆q(calorimeter). ∆q = sm∆T or ∆q = C∆T where C is capacity (mass * specific heat) - Hess's law (meaning: basically you can manipulate chemical equations to achieve a certain reaction and in the process the ΔH value is respectively changed in the process) - bond energies Remember that bond energies are ALWAYS reactants minus products, not to be confused with enthalpy of formation which is ALWAYS products minus reactants. (HBondsBroken-HBondsFormed) - formation enthalpies ΔHo = (ΔHf - ΔHI)

Chapter Eleven - PV - nRT

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R is 62.363 for torr/mmHq, 0.08206 for atm, and 8.314 for Pa. partial pressures and mole fractions Dalton’s Law of Partial Pressures: o PTotal = P1+P2+P3+….PN n gas1 o Mole Fraction (denoted by Χ) = ngas 1+ ngas 2+ ngasn o P1 = X1 (PTotal)

Chapter Twelve - application of IMFs to solids, liquids, and gasses. Be mindful that VAPOR PRESSURE is the ONLY property that decreases as the IMFs increase. Vapor pressure is defined as the ease of breaking an atom free from the collective group, meaning that as IMFs increase, it becomes more difficult to remove the atoms from the group.

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