Chem100 Summary PDF

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Full course summary, focusing on areas highlighted by past exams and lecturer hints...

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Chem100 Summary Measurements and calculations  Significant figures  1.) Non-zero integers are always significant  2.) Leading zeroes are never significant (regardless of decimal point  3.) Captive zeroes are always significant  4.) Trailing zeros are significant, only if they occur after a decimal point. If before decimal point, not significant  For multiplication/division, the number of sig figs in the result is the smallest number of sig figs in the equation  For addition/subtraction, sig figs in the result equals the smallest number of dec places in equation  Atoms  Nucleus is positively charged, contains protons (+ve) and neutrons (no charge)  In a neutral atom, number of protons = number of electrons  Electron configurations  4 principle energy levels  Broken up into S, P, D, F  1 only has S  2 has S, P  3 has S, P, D  4 has S, P, D, F  S=1 orbital, P=3 orbital, D=5 orbital, F= 7 orbital  Each orbital can take a maximum of 2 electrons. If carrying 2, must spin in opposite direction  When get to 3D, 4S is filled up first, and then 3D (and other consequent D levels)  Group (column), skipping transitional elements) = number of valence electrons  Period (row, heading down) gives highest occupied principle energy level. Except for transitional which are one unit behind

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 Noble gas, find element and note its atomic number, and go back to last noble gas that was passed (atomic number)  Relative atomic size: Increases size heading down periods, decreases size heading across group Isotopes  Meaning an atom with a different number of neutrons in its nucleus.  Mass of protons and neutrons are equal.  Can work out number of neutrons by taking the isotope atomic mass, and minus the atomic mass (N0 of protons) Diatomic elements  Outside of compounds, only exist in pair of 2 atoms. I.e. O2.  N to I, 7 shape, plus hydrogen Cations, +ve, Anions, -ve (ide) Naming Compounds  Anion first, then cation  Binary ionic compounds, just name cation first, then anion with an ide ending  Binary ionic compounds with a polyvalent metal, work out how much charge is needed to create an overall charge of zero, name the polyvalent element, follow with charge needed in roman numerals in brackets, followed by anion with ide ending  Compounds that contain only nonmetals, use name suffixes. If first element is only mono, don’t say, just leave as is.  Mono  Di  Tri  Tetra  Penta  Hexa  Hepta  Octa  For polyatomic ions, name first compound, then name polyatomic ion from chart

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 If H appears first, dealing with an acid. If ion doesn’t contain polyatomic ion or oxygen, write hydro xic acid. If It does, Name the polyatomic ion, replace ate with ic and ite with ous, and write acid. Ionic compounds in water  When Ionic compounds are dissolved in water, +ve and -ve ions separate. Creates an electric current. Weak or strong electrolyte Precipitates  Solid  Use punnet square using appropriate charges Equation Types  Molecular equation: Basic formula. Shows states. Doesn’t show separated ions. Remember compounds = charge  Complete ionic equation: Shows everything including charges. Separated ions. States  Net ionic equation: Reactants show involved ions in separated state, showing charges and states. Product shows the two ions combined to create a precipitate Acids  Strong acid completely disassociates in water to H+ ion and anion  Reaction of a strong acid and a strong base produces water  Common strong acids include: HCl, HNO3, H2SO4 Moles  The amount of an element that contains as many atoms as there are in 1g of carbon-12 isotope  1 mole of anything = Avogadro’s number atoms of that thing Percentage Composition  Find molar mass of compound, then the molar mass of element in question, being wary of how may there are. Then dived by molar mass of compound and x by 100. Empirical formula  Simplest whole number ratio of atoms

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 1.) Assume you have 100g and obtain the mass of each element present  2.) Find N0 of moles of each element  3.) Divide all numbers of moles by the smallest mole number, to make at least one of them 1.  4.) Multiply all number of moles by the smallest integer that will make all amounts an integer (or close enough)  To get back to molecular formula, divide molar mass of the compound by the molar mass of the empirical formula and multiply all subscripts in empirical formula by that number. To find element mass from moles, n0 of moles x molar mass Limiting reactants  1.) Make sure equation is balanced  2.) Convert known mass in to moles and re-arrange stoichiometry to allow a 1:? Ratio.  3.) Find out the mass of element in question needed and see if it is less or more than what you have available.  4.) Solve the question based on the limiting reactant by arranging the stoichiometry appropriately and converting the desired product into desired units. Types of bonds  Covalent bond: Sharing on valence electrons such as Ionic bond: Metal and non-metal. Electron transfer  Polar covalent bond: Difference in electro negativity creating a partial positive and partial negative charge.  Bigger the difference, stronger the bond. Electronegativity increase across groups, and decreases down periods  Called a dipole dipole moment Lewis structures  1.) Count the total number of valence electrons in compound/ion and consider any charge on ions.  2.) Connect each bonded atom with 1 line denoting 2

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 3.) Fill out electrons around outer atoms until they reach a count of 8 each or are all used up, including the bond electrons. (8 rule).  If all electrons used up, then problem is finished. I.e., H only needs 2 to fill valance level.  4.) Put any leftover electrons onto central atom. Check to see if totals 8 including bonds. If not, adjust by changing one of the bonds to double or triple bonds to keep 8 in the centre and 8 on outer atoms (inclusive of bonds) Determining electron geometry  Looking at the centres of electrons, doesn’t consider lone pairs of electrons  2 electron groups: Linear, 180 degrees  3 electron groups: Trigonal planer, 120 degrees  4 electron groups: Tetrahedral, 109.5 degrees  If Lewis structure doesn’t have any lone pairs, then the molecular geometry is the same Determining molecular geometry  Consider electron centres and lone pairs  Electron groups will be as far apart as possible.  Shapes will be linear, bent/v-shaped, trigonal planar, trigonal pyramidal, and tetrahedral Conversion factor  Units you want to convert to go on numerator, units wanting to cancel on denominator  Total pressure = sum of partial pressures Gas  Small amounts of gas exert a large amount of pressure, vice versa Ideal gas law  P = pressure in atm  V = volume in litres  n = N0 of moles  R = 0.08206

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 T = temp in kelvin  STP = standard temperature and pressure, Intermolecular forces  The forces that hold different molecules together  Stronger the bonds, the more extreme temperatures required to break bonds to create liquid or vapour.  Solid to liquid energy less than liquid to vapour.  Dipole-dipole  Hydrogen bond is a type of dipole bond which is much stronger than normal. Occurs only with:  H-N  H-O  H-F  London dispersal forces  Occurs in all molecules  Mainly in non-polar molecules  When there is an instant of dipole moment due to electrons orbiting unevenly, which can transfer throughout solution  Weak bond Intramolecular forces  Forces that hold individual atoms together within a molecule  Made up of ionic and covalent bonds  Much stronger than intermolecular forces Types of crystalline solids  Ionic solids  Ions at lattice points i.e. NaCl  Molecular solids  Covalent molecules at lattice points I.e. H2O  Atomic solids  Individual atom at lattice points i.e. C  Bonding in metals, held together by ‘sea of electrons’. Alloys can be made by intisitual metal atom placement into metallic holes Molarity  Molar of solute = n0 of moles/volume in litres

 Ions break up into individual cations. Can calculate molarity. Same as the molarity of solution, but be wary of how many atoms there are of each when broken up.  M1 x V1 = M2 x V2  Redox reactions  OILRIG  Oxidation state rules  1.) Charge/oxidation state of an uncombined element (including diatomic elements) = 0  2.) Charge of monoatomic ion = -1 i.e. Cl 3.) Oxygen is -2 in compounds  4.) Hydrogen is +1 in compounds with non-metals  5.) Sum of oxidation state in compounds = 0  6.) Sum of oxidation states in polyatomic ions = charge on ion  Reducing agent: is itself oxidised  Oxidising agent: is itself reduced  Oxidation: Increase in oxidation number, loss of electrons, reducing agent  Reduction: Decrease in oxidation number, gain of electrons, oxidising agent  Balancing redox reaction  1.) For each half reaction, balance all elements except H and O  2.) Balance O using H20  3.) Balance H using H+  4.) balance charge using electrons  5.) Multiply 1 or both completed half reactions so that the number of added electrons are equal  6.) Add both half reactions and cancel identical species (by charge, i.e. cancel added electrons)  7.) Check elements and charges are balanced  Acids and bases  Acids produce H+ (hydronium) ions, bases produce OH- (hydroxide) ions  Acids are proton/H+ donors and bases are proton acceptors

 Conjugant base is everything that remains of the acid after the H+ proton is lost  Conjugant acid is formed when the proton is transferred to the base  When finding out if 2 elements are a conjugant acid/base pair, they must only differ by one proton. Can’t differ by more than that or a whole A value (rest of acid)  HA + H2O  H3O+ + A A = rest of acid  H2O = base  H3O = hydronium ion, conjugant acid  A- = conjugant base  Weak acid: Most of the acid molecules remain intact/ions don’t separate in base  Has a strong base  Acetic acid CH3OOH  Hydrocyanic acid HCN  Strong acid: almost completely disassociates into H+ and A- ions  Has a weak base  Sulfuric acid H2SO4  Hydrochloric acid HCl  Nitric acid HNO3  pH  H+ = OH- is a neutral solution  H+ > OH- is an acidic solution  OH- > H+ is a basic solution  WITH LOG Ph, the number of decimal places in the answer is equal to the number of significant figures in the original number  pH = 7 is neutral. Under is acidic, over is basic  Buffered solution  Resists change to pH when either an acid or a base has been added  Presence of a weak acid and its strong conjugant base buffers a solution...