Chemistry 101 1st Ediiton Exam 3 Study Guide Chapters 5 8 Lecture 1 Chapter 5 Naming Molecular Compounds Name Molecular Compounds Note that these are molecular compounds with covalent bonds and the ide does not imply an anion in this case In the first element if the first prefix is mono it is usually omitted Lecture 2 Chapter 5 Molar Mass of Compounds Percent Contribution and Chemical Formula Determination Apply the concept of the mole to molecules calculate the molar mass of compounds calculate element percent composition calculate an empirical formula from experimental data Recall that a mole is a number for counting like a dozen or a pair and it is to count atoms The number is 6 02 x 1023 The Molar Mass of a compound is equal to the sums of atomic masses of each individual element within the compound o Ex NaCl molar mass is Na atomic mass 22 99 g mol Cl atomic mass 35 45 g mol 22 99 g mol 35 45 g mol 58 44 g mol of NaCl Its easy to convert between mass moles and molar mass Just remember o Mass of Moles x Molar Mass o Molar Mass Mass of Moles o of Moles Mass Molar Mass practice Mass of 2 5 mol of NaCl Molar Mass of CH4 How many formula units are there in 87 66 g of NaCl Chemical Formula The Empirical Formula gives the relative number of atoms o Think of it as the atom ratio The Molecular Formula gives the actual number of atoms Percent Composition Chemical formulas give the element ration within a compound o E g C6H12O6 has a ration of 1 C 2 H 1 O However each atom has a DIFFERENT atomic mass Therefore the mass ration of element is DIFFERENT than atom ration So How can you calculate percent composition o Step 1 find the mass of a given element in a compound o Step 2 find the molar mass of the compound o Step 3 divide the mass of the element by the mass of the compound o Step 4 multiply by 100 Calculating Empirical Formula from percent composition 1 Convert the percentages to grams a Assume you start with 100 g of the compound 2 Convert grams to moles 3 Write a pseudoformula using the moles as subscripts 4 Divide all by smallest number of moles a If result is within 0 1 of whole number round to whole number 5 Multiply all mole rations by number to make all whole numbers a If ratio is 5 multiply all by 2 if ratio is 33 or 67 multiply all by 3 if ratio is 25 or 75 multiply all by 4 etc b Skip if already all whole numbers Lecture 3 Chapter 6 electronegativity and Bond Polarity Know trends in the periodic table for electronegativity identify pure covalent and ionic bonds qualitatively describe what a dipole is Polar Covalent Bonding Covalent bonding between unlike atoms results in unequal sharing of the electrons This is called a polar covalent bond Electronegativity The ability of an atom to attract bonding electrons to itself is called electronegativity The larger the difference in electronegativity the more polar the bond Bond Polarity Most bonds have some degree of sharing and some degree of ion formation Bonds are classified as covalent if the amount of electron transfer is insufficient for the material to display the classic properties of ionic compounds If the sharing is unequal enough to produce a dipole in the bond the bond is classified as polar covalent Non Polar Covalent Polar Covalent Ionic Bond Bond Dipole Moments dipole moment is a measure of bond polarity generally the more electrons the two atoms share and the larger the atoms are the larger the dipole moment Lecture 4 Chapter 6 Lewis Structures Resonance and Formal Charge Draw Lewis dot diagrams for complex molecules assign formal charges to atoms within molecules construct resonance structures and determine the best resonance structures Lewis structures of Molecules Lewis theory allows us to predict the distribution of valence electrons in a molecule Useful for understanding the bonding in many compounds Allows us to predict shape of molecules Allows us to predict properties of molecules and how they will interact together Exceptions to the octet rule Lewis theory predicts that atoms will be most stable when they have their octet of valence electrons Some atoms commonly violate the octet rule o Ex PCl5 Generally try to follow common bonding patterns Structures that result in bonding patterns different from the common may have formal charges Steps to writing Lewis Structures 1 Draw the correct skeleton structure for the molecule Hydrogen atoms are always terminal The Least electronegative element is usually the terminal 2 Calculate the total number of valence electrons Add an extra electron for each charges and subtract an electron for each charges 3 Place two bonding electrons between each bonded atom Then try to fill the valence shell of atoms by adding electron pairs Start with terminal atom first 4 If any atoms lack an octet form double or triple bonds Do this by moving electron pairs from terminal atoms into the bonding regions 5 Check your work total number of electrons added to the Lewis structure should equal number of electrons in step 2 Formal Charge o During bonding atoms may end with more or fewer electrons than the valence electrons they brought in order to fulfill octets o The result is atoms having formal charge o Formal charge valence electrons nonbonding electrons bonding electrons Resonance o Extensions of Lewis theory suggest that there is some degree of delocalization of the electrons we call this concept resonance o Delocalization of charge helps to stabilize the molecule o Resonance structures differ only in the positions of electrons o The actual molecule is a combination of the resonance forms a resonance hybrid Rules of Resonance Structures Neutral atoms must have a formal charge of 0 Sum of all formal charge must equal the charge of the compound Small formal charges are better than large formal charges Negative charge should be on the most electronegative charge Better structures have fewer formal charges Better structures have smaller formal charges Better structures have the negative formal charge on the more electronegative Steps to drawing resonance structures 1 Draw first Lewis structure that maximizes octets 2 Assign formal charges 3 Move electron pairs from atoms with formal charge toward atoms with formal charge 4 If formal charge atom is 2nd row only move in electrons if you can move out electron pairs from multiple bonds 5 If formal charge atoms is 3rd row or below keep bringing in electron pairs to reduce the formal charge even if get expanded octet Lecture 5 chapter 6 VSEPR