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Chapter 6

Chapter 6 . Chemical Bonding. Chapter Sections:. Introduction to chemical bonding Covalent bonding & molecular compounds Ionic bonding & ionic compounds Metallic bonding Molecular geometry. Section 1:. Introduction to chemical bonding. Introduction to chemical bonding.

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Chapter 6

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  1. Chapter 6 Chemical Bonding

  2. Chapter Sections: • Introduction to chemical bonding • Covalent bonding & molecular compounds • Ionic bonding & ionic compounds • Metallic bonding • Molecular geometry

  3. Section 1: • Introduction to chemical bonding

  4. Introduction to chemical bonding • What is a chemical bond??? A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

  5. Introduction to chemical bonding • Why do atoms bond? They are working to achieve more stable arrangements where the bonded atoms will have lower potential energy than they do when existing as individual atoms.

  6. Introduction to chemical bonding • Types of Chemical Bonding: 1. Ionic – an electrical attraction that forms between cations (+) and anions (-) 2. Covalent – are formed when electrons are shared between atoms 3. Metallic – formed by many atoms sharing many electrons

  7. Introduction to chemical bonding • However…. • Bonds are never purely covalent or purely ionic. • The degree of ionic-ness or covalent-ness depends on property of electronegativity.

  8. Degree of Ionic/Covalent Character in Chemical Bonds 100% 50% 5% 0% Ionic Polar-Covalent Nonpolar-Covalent

  9. Introduction to chemical bonding • Recall what electronegativity is: The degree of attraction that an atom has to electrons that are within a bonded compound. (see page 161)

  10. Introduction to chemical bonding • To determine the degree of ionic-ness or covalent-ness you must take each of the electronegativities for the elements in the compound and subtract them.

  11. Introduction to chemical bonding • If difference is 0-0.3 = nonpolar covalent • If difference is 0.3 – 1.7 = polar covalent • 1.7 and above = Ionic

  12. Ionic/Covalent Character Due to Electronegativity Differences 3.3 1.7 0.3 0 100% 50% 5% 0% Ionic Polar-Covalent Nonpolar-Covalent

  13. Introduction to chemical bonding 2.5 - 2.1 = 0.4 Polar Covalent 2.5 - 0.7 = 1.8 Ionic 2.5 – 3.0 = 0.5 Polar Covalent • Sulfur + Hydrogen • Sulfur + Cesium • Sulfur + Chlorine

  14. Introduction to chemical bonding • In general however… If bonding elements are on opposite sides of the periodic table then they tend to be ionic. If elements are close together, then they tend to be covalent.

  15. Assignment: • Page 177 #3, 4, & 5 • Page 209 #6

  16. Section 2: • Covalent Bonding & Molecular Compounds

  17. Covalent Bonding • What is a molecule? A neutral group of atoms that are held together by covalent bonds. • May be different atoms such as H2O or C6H12O6 • May be the same atoms such as O2

  18. Covalent Bonding • Molecular compounds are made of molecules ….. Not ions! • We represent molecular compounds by chemical formulas that show numbers of atoms of each kind of element in the compound. CH4 - methane

  19. Covalent Bonding • Diatomic molecules are those elements that exist in pairs of like atoms that are bonded together. • There are 7 diatomic molecules: H2 N2 O2 F2 Cl2 I2 Br2

  20. Covalent Bonding Formation of a covalent bond: • When atoms are far apart they do not attract – potential energy is zero. • As they come closer the electrons are attracted to protons but electrons and electrons repel – but e- to p attraction is stronger!

  21. Covalent Bonding • The electron clouds of the bonded atoms are overlapped and form a “bond length.”

  22. Covalent Bonding • Energy is released when these atoms join together with a bond. • Energy must be added to separatethese atoms – called bond energies. • Bond energy is expressed in kilojoules per mole.

  23. Covalent Bonding • Octet Rule – Atoms will either gain, lose, or share electrons so that their outer energy levels will contain eight electrons (H is an exception since it can only have 2 in the outer level). • These electrons that are being gained, lost, or shared are represented by using the electron dot diagrams.

  24. Examples of electron dot notations • 1 valence electron • 3 valence electrons • 5 valence electrons • 7 valance electrons X X X X

  25. Covalent Bonding • Shared electron pairs and unshared pairs: Cl:Cl Shared pair Unshared pairs

  26. Covalent Bonding • These electron dot representations are called Lewis structures. • Dots represent the valence electrons

  27. Lewis structures

  28. Covalent Bonding • Lewis structures can also be represented using structural formulas. • Dashes indicate bonds of shared electrons (unshared e- are not shown Cl - Cl • One pair (2 e-) is shared here.

  29. Covalent Bonding • Lewis structure for ammonia (NH3)

  30. Covalent Bonding • Practice: • Draw Lewis structure for methane CH4 • Ammonia NH3 • Hydrogen Sulfide H2S • Phosphorus trifluoride PF3

  31. Covalent Bonding • Some atoms can form multiple bonds – especially C, O, & N. • Double bonds are bonds that share 2 pair of electrons C=C means C::C • Triple bonds share 3 pair C≡C means C:::C

  32. Covalent Bonding • Resonance: • Some substances cannot be drawn correctly with Lewis structure diagrams • Some electrons share time with other atoms – ex. Ozone – O3

  33. Covalent Bonding • Electrons in ozone may be represented as: O = O–O • Other times it may be represented as O–O=O • Actually these structures are shared – electrons “resonate” (go back & forth) between them

  34. Covalent Bonding • Assignment: p. 189 #4 a – e

  35. Section 3: • Ionic Bonding and Ionic Compounds

  36. Section 3: Ionic Bonding & Compounds • Ionic compounds are formed of positive and negative ions • When combined these charges equal zero Ex: Na = 1+ Cl = 1- 0 charge

  37. Section 3: Ionic Bonding & Compounds • Ionic substances are usually solids • Ionic solids are generally crystalline in shape • An ionic compound is a 3-D network of + and – ions that are attracted to each other

  38. Section 3: Ionic Bonding & Compounds • Crystals in ionic compounds exist in orderly arrangements known as a crystal lattice.

  39. Section 3: Ionic Bonding & Compounds • Ionic substances are not referred to as “molecules” • Ionic substances are referred to as “formula units” • A formula unit is the simplest ratio of the ions that are bonded together.

  40. Section 3: Ionic Bonding & Compounds • The ratio of ions depends on the charges. • What would result when F-combines with Ca2+? • CaF2

  41. Section 3: Ionic Bonding & Compounds • When ions are written using electron dot structures the dots are written and symbols for their charges. • Na.  Na+ • Cl  -

  42. Compared to molecular compounds, ionic compounds: • Have very strong attractions • Are hard, but brittle • Have higher melting points and boiling points • When dissolved or in the molten state they will conduct electricity

  43. Polyatomic Ions: • A group of atoms covalently bonded together but with a charge. • Sulfate SO42- • Carbonate CO32- • Nitrate NO3- • Ammonium NH4+

  44. Section 4: • Metallic Bonding

  45. Metallic Bonding • Metals are excellent electrical conductors in the solid state. • This is due to highly mobile valence electrons that travel from atom to atom. e-

  46. Metallic Bonding • Generally metals have either 1 or 2 s electrons • p orbitals are vacant • Many are filling in the d level • Electrons become delocalized and move between atoms

  47. Metallic Bonding • A metallic bond is the mutual sharing of many electrons among many atoms. • Electrons travel in what is known as the zone of conduction.

  48. Metallic Properties • High electrical conductivity • High thermal conductivity • High luster • Malleable (can be hammered or pressed into shape) • Ductile (capable of being drawn or extruded through small openings to produce a wire)

  49. Metallic Bond Strength • Varies with nuclear charge and number of electrons shared. • High bond strengths result in high heats of vaporization (when metals are changed into gaseous phase)

  50. Section 5: • Molecular Geometry

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