Study Guide 4 Chemistry For Citizens
Study Guide 4 Chemistry For Citizens CHEM 1050 - 01
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This 40 page Study Guide was uploaded by Zaynah allen on Sunday April 3, 2016. The Study Guide belongs to CHEM 1050 - 01 at Georgia State University taught by Dr. S Finnegan in Spring 2016. Since its upload, it has received 15 views. For similar materials see Chemistry for Citizens in Chemistry at Georgia State University.
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Date Created: 04/03/16
Chemical bond Chemical bond: Interaction between two atoms. Lewis theory: 1. Valence (outer most) electrons participate in bonding. 2. Electrons can be transferred. 3. Electrons can be shared. 4. Transfer or sharing electrons occurs in such a way that each atom acquires a stable electronic configuration. This is usually the noble gas configuration or OCTET (except hydrogen, which tries to acquire 2 electrons in the outer most shell). Octet Rule • When atoms bond, they tend to gain, lose, or share electrons to result in eight valence electrons – noble gas configuration • Many exceptions – H, Li, Be, B attain an electron configuration like He • He = two valence electrons, a duet • Li loses its one valence electron • H shares or gains one electron o though it commonly loses its one electron to become H + 2+ • Be loses two electrons to become Be o though it commonly shares its two electrons in covalent bonds, resulting in four valence electrons • B loses three electrons to become B 3+ o though it commonly shares its three electrons in covalent bonds, resulting in six valence electrons – expanded octets for elements in Period 3 or below 3 Lewis symbols And we use dots around the symbol to represent valence electrons pair first two dots for the s orbital electrons put one dot on each open side for first three p electrons then pair rest of dots for the remaining p electrons Problem: Write the Lewis symbols for P and N? (F and F ) - Ionic bonds Positive and negative ions form an ionic bond by electrostatic force of attraction. Lewis symbols helps us predict ionic formulas. (potassium bromide, sodium chloride and aluminum chloride) KBr NaCl AlCl3 Covalent bonds Some atoms prefer to share electrons with other atoms to form covalent/molecular compounds. Cl Cl Cl Cl The shared pair of electrons (bond pair) belong to both the atoms. Lone pair of electrons are also present. Coordinate covalent bonds Both the shared electrons are provided by the same atom. Example formation of hydronium ion (H O3).+ H (from acid) exists as hydronium ion in water. The lone pair of electrons on oxygen are shared with hydrogen. Metallic Bonds • The relatively low ionization energy of metals allows them to lose electrons easily • The simplest theory of metallic bonding involves the metal atoms releasing their valence electrons to be shared as a pool by all the atoms/ions in the metal – an organization of metal cation islands in a sea of electrons – electrons delocalized throughout the metal structure • Bonding results from attraction of cation for the delocalized electrons Chemical bonds might have features of both types of bonds (ionic and covalent) Electrons in a bond are not shared equally in a polar covalent bond. Electrons are pulled towards the more non-metallic element. http://cwx.prenhall.com/petrucci/medialib/media_portfolio/10.html Electronegativity The ability of the atom to compete for electrons with other atoms to which it is bonded. decreases Increases Lower the EN more the metallic character of the element. Higher the EN more the non-metallic character of the element. Polar character in a covalent bond Dipole of a bond Electronegativity difference of the two atoms: Very small value covalent Large value ionic Intermediate value polar covalent Decide on the basis of metallic and non-metallic character. General rules: metal and non-metal Ionic non-metal and non-metal covalent Show dipole in water, carbon dioxide, COS Types of bonds-Summary Ionic bonds Covalent bonds Polar covalent bond http://cost.georgiasouthern.edu/chemistry/general/molecule/polar.htm Covalent bonds – bond energy • In general, the more electrons two atoms share, the stronger the covalent bond – must be comparing bonds between like atoms – C≡C (837 kJ) > C=C (611 kJ) > C−C (347 kJ) – C≡N (891 kJ) > C=N (615 kJ) > C−N (305 kJ) • In general, the shorter the covalent bond, the stronger the bond – must be comparing similar types of bonds – Br−F (237 kJ) > Br−Cl (218 kJ) > Br−Br (193 kJ) – bonds get weaker down the column – bonds get stronger across the period – WHY? 13 Trends in Bond Lengths • Generally bond length decreases from left to right across period – C−C (154 pm) > C−N (147 pm) > C−O (143 pm) • Generally bond length increases down the column – F−F (144 pm) > Cl−Cl (198 pm) > Br−Br (228 pm) 14 Skeletal structures 1. Central atom. 2. Terminal atom. 3. H atoms are terminal atoms. 4. Atoms with less EN values are central atoms. 5. Carbon atoms are always central atoms. 6. Molecules and poly atomic ions generally have compact, symmetrical structures. Lewis structures Strategy: 1. Determine the total number of electrons that must appear in the structure. 2. Ions: Add/subtract electrons. 3. Identify the central atom and terminal atoms. 4. Skeletal structure determination. 5. Every bond in the skeletal structure = 2 electrons. 6. With the remaining electrons first complete the octets of terminal atoms. 7. Then complete the octet of central atom(s). 8. If octet could not be completed then explore the possibility of multiple bonds. 9. If there are more than one possible structures then, check formal charges. Practice – Draw Lewis Structures of the Following CO H PO 2 3 4 2− SeOF 2 SO 3 − NO 2 P 2 4 17 Practice – Lewis Structures 16 e− CO 2 H 3O 4 − 32 e SeOF 2 SO 32− − − 26 e 26 e − NO 2 P 2 4 18 e− 14 e− 18 Formal charge Formal charge arises when atoms do not contribute equal number of electrons to the covalent bond. Formal charge is calculated for individual atoms in the Lewis structure. Formal charge = # of valence electrons in the free atom - # of electrons in lone pairs on the atom - 1/2 (electrons in bond pairs) Practice – Assign formal charges CO H PO 2 3 4 2− SeOF 2 SO 3 − NO 2 P 2 4 20 Practice - Assign formal charges CO 2 H 3O 4 all 0 rest 0 SeOF 2 SO 32− S = +1 Se = +1 − NO 2 P 2 4 all 0 21 Is your Lewis structure correct based on formal charge ? Formal charges arise on certain atoms when atoms do not contribute equal number of electrons to covalent bonds. The sum of the formal charges = charge of the ion Neutral molecule : total formal charge = 0 Formal charge should be as small as possible. (the sum of the absolute values of fc) When formal charge cannot be avoided, negative formal charge should reside on the most electronegatve atom. Practice – Are these the best structures? CO 2 H 3O 4 all 0 rest 0 SeOF 2 SO 32− S = +1 Se = +1 − NO 2 P 2 4 all 0 23 Resonance • Lewis theory localizes the electrons between the atoms that are bonding together • Extensions of Lewis theory suggest that there is some degree of delocalization of the electrons – we call this concept resonance • Delocalization of charge helps to stabilize the molecule 24 Resonance Structures • When there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures • hybridtual molecule is a combination of the resonance forms – a resonance draw it that wayes not resonate between the two forms, though we often • Look for multiple bonds or lone pairs . .. . .. . . .. .. . . O. .S. .. ... . S . O 25 Resonance What are the possible Lewis structures for ozone3O ? Experimental O O bond length: 147.5 pm. Experimental O O bond length: 120.74 pm. Experimental bond length of both the O O bonds in ozone is 127.8 pm. Original structure is a resonance hybrid of both the structures. Drawing Resonance Structures −1 1. Draw first Lewis structure that maximizes octets 2. Assign formal charges −1 3. Move electron pairs from atoms with (−) formal charge toward atoms with (+) +1 formal charge 4. If (+) fc atom 2drow, only move in electrons if you can move out electron pairs from multiple bond rd −1 5. If (+) fc atom 3 row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet −1 +1 27 Drawing Resonance Structures −1 1. Draw first Lewis structure that maximizes +2 octets 2. Assign formal charges 3. Move electron pairs from atoms with (−) formal charge toward atoms with (+) − 1 formal charge 4. If (+) fc atom 2 row, only move in electrons if you can move out electron pairs from multiple bond 5. If (+) fc atom 3 row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet 28 Exceptions to Lewis theory 1. Odd electron species. 2. Incomplete octets (B3 ). 3 Ionic 1 2 Shapes of molecules Shape is the deciding factor for many properties of the compounds. Bond length: distance between the nuclei of bonded atoms. Bond angle: angle between the adjacent lines representing bonds. Important: all diatomic molecules are linear. Valence-Shell Electron-Pair Repulsion (VSEPR) Theory Electron pairs assume orientations around the nucleus to minimize repulsions. This results in particular geometric shape of the molecule. The focus in this theory is not just on pairs of electrons but groups of electrons. Groups of electrons can be bond pairs, lone pairs, electrons involved in multiple bonds and single electrons. Shape of the molecule is determined by orientation of the atoms and not electron pairs. Four electron groups around central atom All of them bonded Write the Lewis structure of CH . How many groups of electrons are surrounding the central atom? http://www.elmhurst.edu/~chm/vchembook/204tetrahedral.html Four electron groups around central atom one lone pair and three bond pairs Write the Lewis structure of NH . How many groups of electrons are surrounding the central atom? Molecular geometry: Trigonal pyramidal http://cwx.prenhall.com/petrucci/medialib/media_portfolio/10.html http://www.uyseg.org/greener_industry/pages/ammonia/1AmmoniaAPQ.htm http://sun.menloschool.org/~dspence/arda/chem_project/web_huey/nitrogen-potassium.html Four electron groups around central atom two lone pairs and two bond pairs Write the Lewis structure 2f H O. How many groups of electrons are surrounding the central atom? Molecular geometry: Angular Memorize table 10.1 http://cwx.prenhall.com/petrucci/medialib/media_portfolio/index.html http://cwx.prenhall.com/petrucci/medialib/media_portfolio/10.html Predicting the shape of a molecule 1. Draw the possible Lewis structure of the molecule. 2. Determine the number of electron groups and distinguish bonded and pairs of electrons. 3. Establish the electron group geometry and then the molecular geometry from table 10.1. Problem: Predict the molecular geometry of nitrogen trichloride. The next important factor to remember: 1. Closer the electron groups, stronger the repulsion between them. 2. The order of repulsive forces Lone pair – lone pair > Lone pair – bond pair > bond pair – bond pair 3. Count multiple bonds of a central atom as one group. 4. A molecule with more than one central atom might have different shapes in different parts of the molecule. Not Quite Perfect Geometry Because the bonds and atom sizes are not identical in formaldehyde, the observed angles are slightly different from ideal 35 The Effect of Lone Pairs • Lone pair groups “occupy more space” on the central atom – because their electron density is exclusively on the central atom rather than shared like bonding electron groups • Relative sizes of repulsive force interactions is Lone Pair – Lone Pair > Lone Pair – Bonding Pair > Bonding Pair – Bonding Pair • This affects the bond angles, making the bonding pair – bonding pair angles smaller than expected 36 Bond Angle Distortion from Lone Pairs 37 Bond Angle Distortion from Lone Pairs 38 Molecular shapes and dipole moments Why is HCl called a polar molecule? As a result of an unsymmetrical distribution of electrons, the bond or molecule contains a positive and a negative pole and is therefore a dipole. The extent of this charge displacement is given by dipole moment, . = d = partial charge, d = distance. Water CO NH 2 3 http://wps.prenhall.com/wps/media/objects/476/488316/ch10.html Predicting polarity of a bond/molecule Understand if you are asked to predict the polarity of a bond or a molecule. A bond will be polar if both the atoms differ in electronegativity. Thus a Cl – Cl , F – F, O – O etc. are non- polar bonds. Whereas, H – Cl, O – H etc are polar bonds. 1. Draw Lewis structure of the molecule. 2. The molecule is polar if the bond dipoles do not cancel out. Some tips: • Planar symmetric molecules are non-polar (Ex: BF ). 3 • Diatomic molecules of identical atoms are non-polar (Ex: O , Cl 2. 2 • If there are no lone pairs on the central atom, and if all the bonds to the central atom are the same, the molecule is non polar. • If the central atom has at least one polar bond and if the groups bonded to the central atom are not all identical, the molecule is probably polar.
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