1 / 44

Unit 12

Unit 12. Chemical Bonding. Definitions. Chemical Bonds Force that holds atoms together It’s all about the electrons (e-) Electrons are attracted to positively charged nucleus of other atom. Types of Chemical Bonds. Ionic Bond

samson-meyers
Download Presentation

Unit 12

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Unit 12 Chemical Bonding

  2. Definitions Chemical Bonds • Force that holds atoms together • It’s all about the electrons (e-) • Electrons are attracted to positively charged nucleus of other atom

  3. Types of Chemical Bonds Ionic Bond • Bond between metal and nonmetal due to “electrostatic interactions” • Attraction between positively and negatively charged ions (cations and anions) • Electrons are transferred from metal to nonmetal

  4. Ionic bonds Result from a Transfer of Valence Electrons + -

  5. Types of Chemical Bonds Covalent Bonds • Bonds in which e- are shared • Most common type

  6. Shared Electrons Complete Shells F F

  7. Types of Chemical Bonds Metallic Bonds • Atoms are bonded to one another (not to other elements) • Positive ions in a “sea” of negative charge (e-)

  8. Metallic Bonding

  9. Definitions • Octet rule (Rule of 8) • 8 e- in the outer shell very stable • H2 and He want a “duet” • Electron configuration for duet = ns2 • Electron configuration for octet = ns2 np6

  10. Examples of Bonding Types • Ionic Bonding: • NaCl, K2S • Covalent Bonding • H2 , Cl2 • Metallic Bonding • Cu, Ag

  11. Lewis Dot Diagrams . . . C . Carbon has four electrons in its valence shell (carbon is in group 14), so we place four dots representing those four valence electrons around the symbol for carbon. • A Lewis dot diagram depicts an atom as its symbol and its valence electrons. • Ex: Carbon

  12. Drawing Lewis Dot Diagrams . . . . . Cl . . • Electrons are placed one at a time in a clockwise manner around the symbol in the north, east, south and west positions, only doubling up if there are five or more valence electrons. • Same group # = Same Lewis Dot structure • Ex. F, Cl, Br, I, At • Example: Chlorine (7 valence electrons b/c it is in group 17)

  13. Paired and Unpaired Electrons • As we can see from the chlorine example, there are six electrons that are paired up and one that is unpaired. • When it comes to bonding, atoms tend to pair up unpaired electrons. • A bond that forms when one atom gives an unpaired electron to another atom is called an ionic bond. • A bond that forms when atoms share unpaired electrons between each other is called a covalent bond.

  14. Writing Lewis Dots Structures for Ions • Uses either 0 or 8 dots, brackets and a superscript charge designate to ionic charge • Ex.) Li+, Be+2, B+3, C+4, N-3, O-2, F-1

  15. Writing Lewis Dots Structures(Ionic Compounds) Lewis Dot Diagrams of Ionic Compounds • Ex. 1) NaCl • Ex. 2) MgF2

  16. Lewis Dot Diagrams for Covalent Compounds • A substance made up of atoms which are held together by covalent bonds is a covalent compound. • They are also called molecules.

  17. Covalent Compounds and Lewis Dot Diagrams • Diagrams show bonds in a covalent compound and tells us how the atoms will combine • Shared e- = bonding e- • Non-shared e- = lone pair e- (a.k.a. non-bonding e-) • Ex. F2

  18. Drawing Electron Dot Diagrams for Molecules • Chemists usually denote a shared pair of electrons as a straight line. F F • Sometimes the nonbonding pair of electrons are left off of the electron dot diagram for a molecule

  19. Examples H CH4 H C H H H N H NH3 H

  20. Types of Covalent Bonds • Single Bond • 2 e- are shared in a bond (1 from each atom) • Double Bond • 2 pairs of e- are shared (4 e- total, 2 from each atom) • Triple Bond • 3 pairs of e- are shared (6 e- total, 3 from each atom)

  21. Rules for Drawing Lewis Dot Diagrams • Add up the total number of valence e- for each atom in the molecule. • Each (-) sign counts as 1 e-, each (+) sign subtracts one e- • Write the symbol for the central atom then use one pair of e- to form bonds between the central atom and the remaining atoms. • Count the number of e- remaining and distribute according to octet rule (or the “duet” rule for hydrogen) • If there are not enough pairs, make sure the most electronegative elements are satisfied. Then, start shifting pairs into double and triple bonds to satisfy the octet rule. • If there are extra e-, stick them on the central atom.

  22. Checking Your Work! • But Remember.... • The Structure MUST Have: the right number of atoms for each element, the right number of electrons, the right overall charge, and 8 electrons around each atom (ideally).

  23. Covalent Compounds and Lewis Dot Diagrams F2 NH3 H2O NH4+ O2 N2

  24. VSEPR: Shapes of Molecules • VSEPR Theory (definition) • “Valence Shell Electron Pair Repulsion” • Based on idea that e- pairs want to be as far apart as possible • Gives molecule its shape

  25. VSEPR: Shapes of Molecules • Electron Pair • Any two valence e- around an atom that repel other e- pairs • Lone pair e- (unshared/non-bonding pair only on one atom) • Shared e- pair (bonding pair shared between two atoms) – can be single, double, or triple bonds

  26. VSEPR: Shapes of Molecules Basic Shapes and Bond Angles

  27. V. VSEPR: Shapes of Molecules

  28. V. VSEPR: Shapes of Molecules

  29. See handout with all of the molecular geometry options!!

  30. To determine the electron pair geometry: • 1. Draw the Lewis structure. • 2. Count the number of bonded (X) atoms and non-bonded or lone pairs (E) around the central atom. • 3. Based on the total of X + E, assign the electron pair geometry. • 4. Multiple bounds count as one bonded atom!

  31. Electron-pair geometry around a central atom Sum of X + EShapes 2 linear 3 trigonal planar 4 tetrahedral 5 trigonal bipyramidal 6 octahedral

  32. VSEPR Examples: What shape would the following compounds have according to VSEPR theory? CH4 CO2

  33. Bond and Molecule Polarity • Polar Bond • Covalent bond in which the electrons are unequally shared • Ex. H2O • Non-polar Bond • Covalent bond in which the electrons are equally shared • Ex. F2 or CH4 • Predicting Bond Polarity • Use Electronegativity!! (see next slide)

  34. Predicting Bond Polarity • Calculate the difference between the Pauling electronegativity values for the 2 elements 0 – 0.4  Non-polar covalent 0.4 – 1.7  Polar covalent (more e/n element has greater pull) 1.7 and up  Ionic (e- are transferred between atoms)

  35. Polar Molecules • Polar Molecules (dipole) • Molecule with separate centers of (+) and (-) charge • In other words, molecules are polar if the pull in any one direction is not balanced out by an equal & opposite pull in the opposite direction

  36. Polar Bonds and Polar Molecules • Drawing Polar Molecules • Positive and Negative regions shown by “delta”(δ) • Ex. CH3Cl

  37. Determining the Polarity of a Molecule • Shape is crucial (determine the VSEPR shape 1st) • All non-polar bonds = nonpolar molecule • Polar bonds  see if they cancel each other out • If they all cancel = nonpolar molecule • If they are unbalanced = polar molecule

  38. Examples: Polar or non-polar? Determine if the following molecules are polar or nonpolar. H2S F2 H2O

  39. Special Types of Bonding Hydrogen Bonding • Force in which a hydrogen atom covalently bonded to a highly electronegative element (F, O, or N) is simultaneouslyattracted to a neighboring nonmetal atom

  40. Hydrogen Bonding • Elements that undergo H-bonding • Hydrogen bonding is FON! • (Fluorine, Oxygen, and Nitrogen) • Effects on Physical Properties • H2O is most notable example of H-bonds • Ice forms rigid, open structures • Increases volume upon freezing (floats) • Molecules w/ higher molar mass have lower BP than H2O

  41. Special Types of Bonding • Van der Waals (London Dispersion) Forces • Intermolecular force between the molecules of a substance • Force of attraction between an instantaneous and induced dipole • Molecules “make these up” (more or less)

  42. Solids • Classes of Solids • Molecular • Formed by molecules containing covalently bonded atoms • Ionic • Formed by cations and anions • Network Covalent • Formed by atoms, usually from Group IV A (Group 14) • Metallic • Formed by positive ions in a “sea” of electrons

  43. Solids Comparison of Solids

More Related