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Chapter 6

Chapter 6 . Chemical Bonding. Section 1:. Introduction to chemical bonding. 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.

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Chapter 6

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  1. Chapter 6 Chemical Bonding

  2. Section 1: • Introduction to chemical bonding

  3. 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

  4. 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- increase stability.

  5. 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

  6. 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.

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

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

  9. 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.

  10. Introduction to chemical bonding • If difference is 0-0.3 = nonpolar covalent • If difference is 0.3 – 1.7 = polar covalent • Above 1.7 = Ionic

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

  12. 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

  13. Introduction to chemical bonding • In general however… If bonding elements are on opposite sides of the periodic table (metal with a nonmetal) then they tend to be ionic. If elements are close together (nonmetal to nonmetal), then they tend to be covalent.

  14. Section 2: • Covalent Bonding & Molecular Compounds

  15. 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

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

  17. 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 • Big 7

  18. 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!

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

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

  21. 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.

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

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

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

  25. 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.

  26. Steps To Drawing Lewis Structures • Calculate the number of valance electrons. • Arrange atoms. • Compare number of electrons used with number of electrons available. • Check octet rule. • Change dots to dashes where appropriate.

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

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

  29. More Guidelines • H and halogen atoms usually bond to only one other atom in a molecule and are usually on the outside or end of a molecule (each only need 1 electron to form stable octet and electronegativity)

  30. More Guidelines • The atom with the smallest electro-negativity is often the central atom • When a molecule contains more atoms of 1 element than the other, these atoms often surround the central atom

  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. Section 3: • Ionic Bonding and Ionic Compounds

  35. 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

  36. 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

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

  38. 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.

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

  40. 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  -

  41. 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

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

  43. Section 4: • Metallic Bonding

  44. 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-

  45. 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 (sea of electrons)

  46. Metallic Bonding • A metallic bond is the mutual sharing of many electrons among many atoms.

  47. 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)

  48. 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)

  49. Section 5: • Molecular Geometry

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