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UNIT 9: Covalent Bonding 1 What is a Covalent Bond? • Covalent Bond –formed when two nonmetals share pairs of valence electrons in order to obtain the electron configuration of a noble gas • Molecule - formed when two or more atoms bond covalently. (A molecule is to a covalent bond as a formula unit is to an ionic bond.) 2 Diatomic Molecules • HOFBrINCl Share electrons when they bond together 3 Polyatomic Ions • covalently bonded group of atoms, with a charge 4 Properties of Covalent Molecules • Can exist as gases, liquids, or solids depending on molecular mass and polarity • Usually have lower MP and BP than ionic compounds of the same mass • Do not usually dissociate (break apart into ions) in water • Do not conduct electricity 5 How to draw Lewis dot structures for covalent molecules: 1. 2. 3. Write the formula for the compound. Count the total number of valence electrons. Predict the location of the atoms: a) If there is only 1 atom of an element, it is the central atom. b) If carbon is present, it is ALWAYS the central atom. c) The least electronegative atom is generally the central atom. d) Hydrogen is NEVER the central atom. 4. Place one electron PAIR between the central atom and each ligand (side atom) to “hook” the atoms together. 5. Dot the remaining electrons in pairs around the compound to complete the octet. Start with the ligands. 6. Check that each atom has an octet. (H only needs a pair, not an octet.) 6 Lewis Structures for Molecules • Draw the Lewis dot structure for these molecules: – Hydrogen + Bromine (HBr) – Carbon + Chlorine (CCl4) 7 Writing Lewis Dot Structures - Covalent Bonds Bonding e- Pairs Lone Pairs (nonbonding electrons) Covalent bonds Exceptions to the octet rule: • Molecules that have an odd # of valence electrons; ex. NO2 has 17 total valence electrons and can’t form an exact # of pairs • Molecules with fewer than 8 electrons present; ex. BH3 where B only has and only needs 6 electrons • Molecules with an expanded octet; ex. PCl5 where P forms 5 bonds and SF6 where S forms 6 bonds 9 Number of bonds • Single Bonds - when one pair of e- is shared between atoms • Double bond – when atoms share 2 pairs of valence electrons; ex. O2 • Triple bond – when atoms share 3 pairs of valence electrons; ex. N2 10 Describing bonds • Sigma bond - the first bond between 2 atoms – A single bond is a sigma bond. • Pi bond - the second bond between 2 atoms – A double bond consists of a sigma bond and a pi bond. – A triple bond consists of a sigma bond and two pi bonds. 11 Carbon can form single, double and triple bonds with itself. 12 Types of Bonds: • Nonpolar covalent (also called pure covalent or covalent) equal sharing of electrons between atoms; occurs between the atoms in a diatomic molecule (HOFBrINCl) and between C and H; ex. CH4 • Polar covalent – unequal sharing of electrons between atoms; occurs between two nonmetals or a nonmetal and a metalloid; ex. H2O • Ionic – complete transfer of electrons; occurs between m/nm, m/PAI, PAI/nm or PAI/PAI; ex. NaCl 13 Bond type • • Difference in electronegativity values Distance between atoms on the periodic table Non-Polar Covalent NPC Small Polar Covalent PC Ionic I medium big THIS IS A CONTINUUM. IT DESCRIBES THE “IONIC CHARACTER” OF THE BOND. 14 Practice: • What type of bond exists in each of the following? • 1. HCl 2. CaO • 3. H2O • 4. Br2 15 Why are molecular shapes important? The shape of a molecule plays a very important role in determining its properties. Properties such as smell, taste, and proper targeting (of drugs) are all the result of molecular shape. Molecular Shape Lewis structures do not show how atoms in a molecule are arranged in 3-dimensional space. Can you tell the molecular shape of CCl4 from its Lewis structures? www.mikeblaber.org Molecular Shape Example: Water is not linear! http://chemistry.tutorvista.com/ Atoms in a molecule try to spread out from one another as much as possible to reduce the charge repulsion between their outer electrons. H methane, CH4 H C H Is this the farthest that the hydrogens can get away from each other? H 90° 109.5° science.howstuffworks.com science.howstuffworks.com This shape causes less repulsion between the bonding pairs of electrons as the hydrogen atoms are farthest away from each other. Molecular Shape apchemcyhs.wikispaces.com Molecules adopt a geometry (shape) that minimizes e – e repulsions. This occurs when e- pairs are as far apart as possible. Sample problem – molecular geometry What is the shape of the following molecules? commons.wikimedia.org http://winter.group.shef.ac.uk/ tetrahedral trigonal planar en.wikipedia.org www.chriscrews.com Bent or angular Trigonal pyramid VSEPR • Valence Shell Electron Pair Repulsion • A theory that states that electron pairs repel both bonding and non-bonding electrons resulting in a stable (lowest-energy) 3dimensional geometry. 22 TO DETERMINE MOLECULAR SHAPE • Use VSEPR (valence shell electron pair repulsion) rules: 1) Draw the Lewis dot structure for the molecule 2) Identify the central atom 3) Count total # of electron pairs around the central atom (stearic number) 4) Count # of bonding pairs of electrons (regions of electron density) around the central atom 5) Count # of lone pairs of electrons around the central atom; lone pairs take up a lot of space 6) Look at summary chart, identify shape **shapes with no lone pairs are symmetrical **shapes with lone pairs are assymmetrical 23 Practice: • Determine the shape. 1. NF3 2. SiCl4 3. H2O 24 Water is a POLAR molecule The more electronegative atom will have a slight negative charge, the area around the least electronegative atom will have a slight positive charge. 25 Symmetric molecules tend to be nonpolar Asymmetric molecules with polar bonds are polar 26 Naming Binary Molecules 1. Write the name of the first element. 2. Change the nonmetal’s ending to –ide. 3. Use prefixes to indicate the number of each type of atom. *Exception-the first element will never have the prefix mono – – – – – – – – – – 1-mono 2-di 3-tri 4-tetra 5-penta 6-hexa 7-hepta 8-octa 9-nona 10-deca 27 Practice • Write the name for the following molecules: 28 Writing formulas for molecules • The prefixes tell you the subscript for each atom. 29 Practice • Write the formulas for the following molecules: 30 Empirical formula • Empirical - To be derived from observation, experiment, or data. • Empirical formula - the simplest whole number ratio between two (or more) elements 31 Steps to determine the empirical formula of a compound: 1. Determine the mass of each element in the sample. 2. Divide the mass of each element by the molar mass (from the PT) to determine the number of moles of each element. Round to the thousandths (._ _ _ )! 3. Divide the # of moles of each element by the smallest # of moles. This is the mole ratio for each element in the compound. 4. If your answers to step 3 are whole numbers, these are written as the subscripts. 5. If your answers to step 3 are NOT whole numbers, multiply by 2, 3, or 5 to obtain a whole number if increments of 0.5, 0.3 or 0.2 are given, respectively. 32 example What is the empirical formula for a sulfur oxide compound containing 50% sulfur and 50% oxygen? Step 1: Since % means “parts per hundred”, assume we are working with a 100 g sample. That means we have 50 g of sulfur and 50 g of oxygen. 33 Step 2: Use dimensional analysis to convert grams to moles 50 g S __1 mol_ = 1.558 moles S 32.1 g 50 g O _1 mol_ 16.0 g Label Properly!! = 3.125 moles O Round to thousandths (._ _ _) 34 Step 3: Divide by the smallest number of moles to obtain a mole ratio. 1.558 moles S =1S 1.558 moles S 3.125 moles O Label Properly!! =2O 1.558 moles S So, we have 1 S for every 2 O. These numbers become the subscripts and the formula is SO2 In our example, we did not need the 4th step since the ratio came out to a whole number. 35 Another example A compound contains 54.1 g of Mg and 45.9 g of P. Determine the compound’s empirical formula. Note: This time, we already have the number of grams so we can skip to step 2. 36 Step 2: Use Dimensional Analysis to convert grams to moles. 54.1 g Mg___1 mol_ = 2.226 moles Mg 24.3 g Label 45.9 g P __1 mol_ = 1.481 moles P Properly!! 31.0 g Round to thousandths place (._ _ _ ) 37 Step 3: Divide by the smallest number of moles to obtain a mole ratio. 1.482 moles P =1P 1.482 moles P 2.226 moles Mg Label Properly!! = 1.5 Mg 1.482 moles P Notice, the bottom answer did not come out to a whole number this time. 38 Skip to step 5 since answers are not whole numbers. Step 5: Multiply answers from step 3 so that you get whole numbers. We had 1.5 Mg and 1 P 0.5 = ½ flip it and you have your scale factor, 2. 1.5 x 2 = 3 Mg and 1 x 2 = 2 P. The ratio did not change, it is just a whole number ratio now. So, we have 3 Mg for every 2 P or the formula Mg3P2 39 Now that you know the steps, here is a jingle to make them easier to remember: Percent to mass step 1 Mass to mole step 2 Divide by small step 3 Multiply ‘til whole steps 4 and 5 Note: You may not need all of the steps. 40 Molecular formula A formula that is reducible. It is a multiple of an empirical formula. Ex. Can C8H12 be reduced? Of course, it’s divisible by 4. So, dividing by 4 reduces the formula to C2H3. C8H12 is the molecular formula. C2H3 is the empirical formula. 41 This template can help you organize your information and find what you are missing. Empirical formula molecular formula Mass of empirical formula Mass of molecular formula 42 Ex. The molar mass of a molecular formula is 283.88 g/mole and it’s empirical formula is P2O5. Determine the molecular formula. Draw your chart and fill in the info from the problem. P2O5 ? 141.943 g/mole 283.88 g/mol Now, divide the molar mass of the MF by the molar mass of the EF. (283.88 g/mole)/(141.943 g/mole) = 2. Scale factor is 2. Multiply the subscripts in the EF by 2 and the MF is… P4O10 43 practice A compound is made from 2.00 g carbon, 0.335 g hydrogen, and 2.66 g oxygen. Its molar mass is 90.0 g/mole. Determine the molecular formula. 44