- Converting Moles (mol.) to grams (g):

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1 Study Guide: Avogadro's # 1mol = 6.02x10^23 particles/atoms/ions/elephants use this as a conversion factor to calculate atoms in compound 1mol C / 6.02 x 10^23 atoms C Percent composition: [Mass of Element x Subscript / Molar Mass of Compound (Total mass)] x 100 = percent composition - Converting Moles (mol.) to grams (g): - Molar mass (g/mol) of a single element is the mass found on the periodic table Empirical Formulas - Empirical Formula: Represents the ratio between the number of each atom found in the compound. The ratio is a mole to mole ratio. Empirical formulas don t always represent the actual structure of the compound. They only tell you the correct ratio. There are 3 possible starting points for an empirical formula calculation: (1) % comp-(2) Grams-(3) Moles-Mole Ratio (subscripts) (1) - (2) : change percent to grams (assume 100 grams) (2) - (3) : divide by molar mass Step 4 (mole ratio): divide all by smallest number of moles Why assume 100g? Makes conversion simple. Why divide by the smallest number of moles? Allows you to make a mole ratio where the smallest number is 1. What if the ratio isn t a whole number? Ex) 1.5 mol O *2 = 3 1molN*2=2 Ends in....1 or.9 5 round up/down.2 or.8 multiply by 5.3 or.7 multiply by 3.4 or.6 multiply by 5.5 multiply by 2 Empirical Formulas of Hydrocarbons(Compounds with only C and H) or Compounds with only C, H, O For Hydrocarbons Problem will most likely give you starting masses of the products of a combustion reaction 1) convert to mol of CO2 and H2O 2) use the # ratio of H in the total molecule(h20) to H ex: 1mol (H20 / 2 mol of H) because the subscript tells you that there are two Hs per H2O 3) do the same for CO2 and solve for C 4) divide by the smallest number and multiply to make both quotients whole numbers For compounds also with O in the formula, they will give you the total mass of the compound 1) Do steps 1-3 from above 2) take those two mol calculations(make sure to keep them though because they will be needed for step 5) and convert to grams using molar mass 3) add those grams together and subtract that from the original total you were given in the question 4) convert this O to mol using the molar mass 5) (step 4) for the moles of the C, H, O

2 - Molecular Compounds: Share the same ratio of atoms, but their actual structures differ. How do we tell the difference between compounds? Represents the actual structure of the compound but shares the same ratio as the empirical formula. Easy way to remember the difference between types of formulas: Empirical: simplest ratio, not structure Molecular: structure and ratio How to convert between empirical and molecular formulas: Molecular Formula Mass / Empirical Formula Mass Multiply subscripts of empirical by the whole number to get the Molecular Formula Excess reactant: 1. convert both starting (g) of reactants to moles of the product (use molar mass + coefficients) 2. take the smaller mol of the product (which ever reactant this comes from is the limiting reactant) and convert to g of the product, This is the theoretical yield Extra: if they ask for the percent yield, then they will give you the actual yield in the problem 3. (Actual yield / Theoretical yield)x100= percent yield higher the percent yield, the more efficient the reaction was because theoretical yield is at 100% efficiency and the closer the actual yield is to the theoretical, the higher the percent yield Balancing Chemical Equations - Conservation of Mass: Mass can t be created or destroyed. - Conservation of mass tells you two important things about equations: 1) The amount of each element must be the same on both sides of the equation. 2) The types of elements must also be the same on both sides of the equation. - We balance reactions to satisfy the Law of Conservation of Mass. - What are you allowed to do to balance an equation? o Change the coefficients of compounds - What are you not allowed to do to balance an equation? o Change the subscripts of an element o Change the elements - TIPS: o Treat polyatomic ions as a unit or group o Write water (H2O) as HOH o No fractions in final answers o Reduce coefficients only if you can reduce all of them o For combustion reactions balance in the following order: H-> C-> O Reaction types: - Combustion (1): o A reaction where a compound, usually a hydrocarbon (mostly made of carbon), reacts with oxygen gas (O2) to form carbon dioxide (CO2) and water (H2O) - Synthesis (2):

3 o A reaction where simple reactants (elements or products) form a single compound. o General form: A+B->AB - Decomposition (3) o A reaction where one complex compound breaks into multiple simple products o Opposite of synthesis o General form: AB->A+B - Single Replacement/ displacement (4): General form: A + BC->AC + B - Double Replacement/ displacement (5): - A reaction where two ionic compounds switch their cations and anions - General form: AB +CD -> AD +CB - Acid Base reaction: type of double replacement - has H in reactant Trimester 2 stuff: Lewis structures show how electrons are exchanged in reactions/bonds o Valence electrons are the only type of electrons that are exchanged (i.e. core stays untouched) - Two types of Lewis diagrams o Dot diagrams (Lewis dot diagrams): one single atom o Lewis structures: two or more atoms - Dot Diagrams o Useful for determining the charge of an ion an atom forms and for determining the number of bonds an atoms likes to form in a covalent structure o Symbol with valence electrons (s and p sublevels only) o Covalent molecules: only single electrons (half filled orbitals) will bond (AKA be shared) Halogens: usually prefers only 1 bond Group 16: usually prefers 2 bonds Group 15: usually prefers 3 bonds Group 14: usually prefers 4 bonds - Drawing Lewis Structures o Steps: 1. Add total number of valence electrons available for bonding 2. Single bond the outer atoms to the central atom 3. Add lone pairs to all atoms to satisfy the octet rule 4. Count electrons and add double or triple bonds if necessary o IMPORTANT: You must always have the same number of electrons in the structure as calculated in step 1! o Tips: H only needs two electrons (not 8) (only 1s) H is always terminal (at the end) More electronegative atoms are terminal Charges on polyatomic ions require you to add or subtract electrons from the total number of electrons Lines= two electrons Exceptions to the Octet Rule o (1) Less than octet: B, and sometimes Be B only needs 6 electrons Be only needs 4 electrons

4 o (2) More than an octet (elements with n=3 period 3 and going down or greater) Result of having d orbitals available for bonding Extra electrons? Look at the central atom, if n=3 or greater, add extras as lone pairs (LPs) to central atom C, N, O, and F NEVER have an expanded octet o (3) Odd number of electrons (odd electron species)(free radicals) **don t really need to worry about this** Very reactive because they are unstable Are able to have a single electron surrounding an atom Resonant Structures o Same compound (chemical formula) but different arrangements of electrons (bond in different places) o Resonance Hybrid: mix of all possible structures o The actual structure of a molecule is an intermediate between the two or more resonance structures (def. of resonance hybrid) Formal Charge o Tells you which resonance structure is the most stable o Charges are assigned to individual atoms in a molecule o Actually fictitious charges (bookkeeping) o FC= # of VEs of the atom - # of owned electrons in structure = #of electrons shared with other atom - # of electrons it ends with o Rules: Sum of all FCs in the molecule must add to equal the total charge of the molecule Neutral = zero Ion = charge Smaller FCs are more stable If you must make an atom have a negative FC, the most electronegative charge prefers it. o Nodes count as 1 o Electrons (LPs) count as 2 Resonance Hybrid: The actual structure of a molecule is an intermediate between two or more resonance structures. o Delocalization: Delocalized electrons are electrons in a molecule, ion or solid metal that are not associated with a single atom or one covalent bond. Delocalized electrons are contained within an orbital that extends over several adjacent atoms. o Polar: Unequal sharing of electrons. o Non-polar: Equal sharing of electrons. o Dipole moment: Being pulled towards the more electronegative atom Partial charge: represented by the Greek lowercase letter δ, namely δ or δ+. Partial charges are created due to the asymmetric distribution of electrons in chemical bonds. Bond Polarity: Electrons are not equally shared in both bonds, between two atoms and molecules. The unequal sharing is due to differences in electronegativity of the individual atoms. To quantify the polarity (how its sharing) of a bond we use a property called dipole moment. Cl2: Equal sharing= same molecule, same electronegativity. Non polar bond HCl: Cl is more electronegative. Polar bond= difference in electronegativity HF: Polar bond- more or less polar than HCl? More polar bond due to a larger change in electronegativity Bond polarity is helpful, but molecules are made up of more than 1 bond

5 - NOTE: BE PREPARED FOR PICTURE PROBLEMS