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W7 Chemical Principles notes

by: Olivia Lange

W7 Chemical Principles notes sch 100 01

Marketplace > Seton Hill University > Chemistry > sch 100 01 > W7 Chemical Principles notes
Olivia Lange
Seton Hill University

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week 7, (10/5, 10/7, 10/12)
Chemical Principles
Professor Flowers
Class Notes
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This 10 page Class Notes was uploaded by Olivia Lange on Sunday October 16, 2016. The Class Notes belongs to sch 100 01 at Seton Hill University taught by Professor Flowers in Fall 2016. Since its upload, it has received 3 views. For similar materials see Chemical Principles in Chemistry at Seton Hill University.


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Date Created: 10/16/16
Chapter 4 Continued... 4.4 Coordinate Covalent Bonds ● A coordinate covalent bond is the covalent bond that forms when both  electrons are donated by the same atom. ● See colorful graphic on page 108 ● Once formed, no different from any other covalent bond ○ Result in unusual bonding patterns ○ Nitrogen with 4 covalent bonds +¿ ○ Oxygen with 3 bonds ( H 3 ¿ ) O ● See graphic at very bottom of page 108 Remember: metals form ionic compounds (giving up), non­metals form covalent (sharing) 4.5 Characteristics of Molecular Compounds ● Ionic compounds­ high melting and high boiling points ○ Attractive forces between oppositely charged ions are so strong. ● But….Molecules are neutral ○ No strong attraction to hold them together. ● And….Weaker forces exist between molecules, called intermolecular forces ● Very weak intermolecular forces = gas ● Somewhat stronger intermolecular forces = liquid ● Strongest = molecular solid ● Molecular solids ○ Lower melting and boiling points than ionic compounds ○ Rarely soluble in water ○ Do not conduct electricity when melted ● See table 4.1 on page 110 ● Molecular formula ○ A formula that shows the numbers and kinds of atoms in one  moleccule of a compound ○ Li ❑ PO  ❑ 3 4 ● Structural formula ○ A molecular representation that shows the connections among  atoms by using lines to represent covalent bonds ○ CBr3   is Br (north, west, east) and C in the middle. All connected  by a single line ● Lewis Structure ○ A molecular representation that shows both the connections  among atoms and the locations of lone­pair valence electrons ● Lone pair ○ A pair of electrons that is not used for bonding See picture 4.7 Drawing Lewis Structures ● Carbon ○ Forms 4 covalent bonds, often bonds to other carbon ● Nitrogen  ○ Forms 3 cov. Bonds, has 1 lone pair ● See figure 4.3 on page 104 ● Oxygen wants 2 cov. Bonds, has 2 lone pair e­ ● Halogens (X) forms 1 cov. Bond, 3 lone pair e­ ● Hydrogen forms 1 cov. Bond ● See page 111 for a visual ● Larger organic molecules are often written as condensed structures in which the  bonds are not specifically shown. ● I.e.: Ethane is  CH 3 CH 3 ● Ball and stick model with pegs and “lego­like” pieces for a visual ● C is black, H is white, O red, N blue, S yellow, P dark blue, etc ● A general method for drawing lewis structure ○ Find the total number of valence electrons of all atoms in the  molecule or ion in  PCl 3 ○ P has 5 valence e­ ○ Cl has 7 ○ 7e­ x3 is 21 e­ plue 5 e­ is 26 e­ ○ Atoms in the 3rd row and beyond of the periodic table occur as the central atom in a cluster ○ In  PCl 3 , you will draw the bond­lines, then add the electron  dots! ○ Place any remaining in central atom (lone pairs) ○ If the central atom does not have an octet after all the electrons  have been assigned, take a lone pair from a neighboring atom and form a  multiple bond to the central atom. Worked Example 4.7: Draw lewis structure for HCN atoms are connected in H, C, N format. Total up your electrons: H gives 1 C gives 4 N gives 5 Total: 10 Now make the bonds… H­C­N Minus 4 e­ because of bondlines… total 6 e­ Draw the 6 electrons around N (two north, two east, two south) Now all electrons are gone! C is unhappy! Borrow two pairs of lone electrons, now C to N is 3 lines and N has two lone on its east side. H ­ C N: ≡ Worked Example 4.8 Lewis Structure for  C 2 H 3 Cl ❑ C gives 4...x2 is 8 H gives 1...x3 is 3 Cl gives 7….. Is 7 Total 18 H    H     now we have 8e­ because the 5 lines (­ and I )take up 2 each  I     I      C ­ C  I     I  H   Cl H    H       I     I      C = C give Cl 8 dots and the middle C bond is now a double (=)  I     I H   :Cl:        ** Worked Example 4.9 SO 2     electrons… S gives 6, O gives 6 x2…. 18 e­ total O ­ S ­ O  gives us 14 e­ total  ..  .. :O ­ S ­ O: gives us 2 e­ total (14­ 12 new dots)  **  **  ..     .. :O ­  S = O:  **    **    done. 4.8 The Shapes of Molecules ● Valence­shell electron­pair repulsion (VSEPR) model ○ Method for predicting molecular shape ○ Constantly­moving valence electrons in bonds and lone pair make up negatively charged clouds of electrons ○ Electron clouds repel one another ○ Clouds tend to keep as far apart as possible ○ This cause molecules to assume specific shapes. ● Applying the VSEPR model ○ STEP 1: draw a lewis structure of the molecule, and identify the  atom whose geometry is of interest­ usually the central atom ○ STEP 2: count the number of electron charge clouds surrounding  the atom of interest. This is the total number of lone pairs plus connections to  other atoms. ○ STEP 3: predict molecular shape by assuming that the charge  clouds orient in space so that they are as far away from one another as possible. ● See table 4.2 in book, the chart will be given with the exam. ● A bond angle is the angle formed by 3 adjacent atoms in a molecule. ● Two charge clouds are farthest apart when they point in opposite directions. ● Linear molecules result with bond angles of 180 degrees. ● See the visual on the bottom of page 117 ● Three charge clouds are farthest apart if they lie in a plane and point to the  corners of an equilateral triangle. ● Trigonal planar molecules result with bond angles of 120 degrees. ● See the top visual at the bottom of page 118  ● If one of the charge clouds is a lone pair, the molecule will be bent. ● See the bottom visual at the bottom of page 118 ● Four charge clouds ○ Extend to the corners of a regular tetrahedron ○ Bond angles of 109.5 degree ● See top visual on page 119 ● Lone pairs repel strongly ● Bond angles reduced slightly ● Depending on the number of lone pairs, shapes can be: ○ Tetrahedral ○ Pyramidal ○ Bent ● See bottom visual on page 119 and top visuals on page 120 Worked Example 4.10 +¿ Shape for hydronium ion, H 3 ¿ O And  CH 3 CHO ❑ 4.9 Polar Covalent Bonds and Electronegativity ● When atoms are not identical, bonding electrons may be shared unequally. ● A polar covalent bond is one in which the electrons are attached more strongly  by one atom than by the other. ● See visual on 122 ● Electronegativity ○ Ability of atom to attract electrons ● Fluorine ○ Most electronegative element ○ Assigned a value of 4 ● Less electronegative atoms are assigned lower values ● See chart on page 123 (Figure 4.6) ● Electronegativities ○ Metallic elements ­ low ○ Halogens and reactive nonmetal elements ­ higher ○ Generally decreases going down the periodic table within a group ● Electronegativity differences ○ Less than 0.5 = nonpolar covalent bond ○  0.5­1.9 = increasingly   polar cova  bonds ○ 2 or more = ionic bonds Worked Example 4.11 C and Br 2.5 ­ 2.8 = 0.3, nonpolar covalent bond Li and Cl 1.0 ­ 3.0 = 2.0, ionic bond N and H 3.0 ­ 2.1 = 0.9, polar covalent, N is partial negative , H is partial positive ¿ ¿ ¿ Si and I 1.8 and 2.5 2.5 ­ 1.8 = 0.7, polar covalent Si is partial positive (S+) and I is partial negative 10/10/16 4.10 Polar Molecules… ● Molecular polarity depends on the shape of the molecule. ● Symmetrical molecules can have polar bonds and be non­polar overall. ● Polarity effects: ○ Physical properties of molecules ○ Example: Melting points, boiling points, solubility 4.11 Naming Binary Molecular Compounds ● Binary compound ­ when two different elements combine ● The less­electronegative element is written first ○ Metals are always written before nonmetals ○ A nonmetal farther left on the periodic table generally comes  before a nonmetal farther right. ● Requires identifying exactly how many atoms of each element are included. ● See table 4.4 on 127 ● “ONLY IN MOLECULAR, NOT IONIC” ● 1.) Name the first element in the formula, using a prefix if needed. ● 2.) Name the second element, using an ­ide ending and a prefix if needed. ● See visual at bottom of 127 ● Prefix mono­, meaning one, is omitted ● Except where needed to distinguish between Worked Example 4.13 Naming Molecular Compounds (numbers are lower subscripts) ● N2O3 is… dinitrogen trioxide ● GeCl4 is…germanium tetrachloride ● PCl5 is… phosphorous pentachloride End of chapter 4 5.1 Chemical Equations ● Chemical Equation ● An expression in which symbols and formulas are used to represent a  chemical reaction. ● “Recipes” ● Starting materials and final products are listed ● Chemical names replaced with formulas 2NaHCO 3  ­>  Na 2 CO 3  +  H 2  + CO2 ● Reactant ● Substance that undergoes change. ● Written on the left side of the reaction arrow. ● Product ● Substance formed. ● Written on right side of the reaction arrow. ● Necessary conditions are written above the arrow. ● Law of conservation of mass ● Matter can neither be created nor destroyed in a chemical reaction. ● Chemical equations must be balanced. ● Bonds between atoms in the reactants are rearranged to form new  compounds. ● None of the atoms disappear. ● No new atoms are formed. ● Numbers and kinds of atoms must be the same on both sides of the  reaction arrow. ● Coefficient ● Number places in front of a formula to balance a chemical equation. ● Coefficients multiply all the atoms in a formula. ● Important to consider coefficients when balancing an equation. ● 3  H2O …. 3 is the coefficient Symbols after the formulas indicate ● (s) = solid ● (l) = liquid ● (g) = gas ● (aq) = dissolved in aqueous solution  Worked Example 5.1 Reactants:........Products Pb: 2 …………. 2 S: 2 ……………. 2 O: 6 …………….. 6


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