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Introduction to Chemical Bonding: Ionic, Covalent, and Metallic Bonds

This chapter provides an introduction to chemical bonding, including the definition of chemical bonds, the types of bonding (ionic, covalent, metallic), and the role of electronegativity differences. It also covers topics such as molecular compounds, Lewis structures, and the properties of ionic and metallic compounds.

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Introduction to Chemical Bonding: Ionic, Covalent, and Metallic Bonds

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  1. Chapter 6 Chemical Bonding

  2. Section1 Introduction to Chemical Bonding Chapter 6 Objectives • Definechemical bond. • Explain why most atoms form chemical bonds. • Describe ionic and covalent bonding. • Explain why most chemical bonding is neither purely ionic nor purely covalent. • Classify bonding type according to electronegativity differences.

  3. Chemical Bond – atoms bound together by a mutual electrical attraction between nuclei and valence electrons • Ionic Bond – bond between anions (-) and cations (+) caused by the transfer of electrons between atoms – usually a metal and a nonmetal • Covalent Bond – bond between atoms that sharevalence electrons – usually 2 nonmetals • Metallic Bonding – attraction between metals with delocalized electrons – (not a real bond) • Hydrogen Bond – attraction that occurs between hydrogen atoms in a compound and strongly electronegative atoms with lone pairs of electrons (oxygen, nitrogen, etc.) (not a real bond) ionic covalent metallic hydrogen Section 1Introduction to Chemical Bonding

  4. Predicting Bond Type by Electronegativity Difference

  5. Section2 Covalent Bonding and Molecular Compounds Chapter 6 Objectives • Definemolecule and molecular formula. • Explain the relationships among potential energy, distance between approaching atoms, bond length, and bond energy. • State the octet rule.

  6. Section2 Covalent Bonding and Molecular Compounds Chapter 6 Objectives, continued • List the six basic steps used in writing Lewis structures. • Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both. • Explain why scientists use resonance structures to represent some molecules.

  7. Polar Covalent Bonding Polar Covalent Bonding results in an unequal sharing or uneven electron distribution Section 2Covalent Bonding and Molecular Compounds

  8. Molecule – • Compound – • Chemical Formula – • Molecular Formula – . • Molecule – group of atoms that are covalently bonded • Compound – group of atoms that are ionicly bonded • Chemical Formula – tells the kind and number of atoms in a compound • Molecular Formula – tells the kind and number of atoms in a molecule

  9. Covalent Bond Characteristics bond length • Bond Length – average distance between two bonded atoms at their minimum potential energy • Bond Angle – angle between three bonded atoms • Bond Energy – energy required to break a chemical bond and form neutral isolated atoms bond angle bond energy

  10. Lewis Structures Structural Formula • Structural Formula – shows the kind, number and arrangement of atoms in a molecule (may or may not show lone pairs) • Octet Rule – compounds tend to form so that both atoms have eight electrons in their outer energy level • Exceptions – hydrogen & expanded octets • Electron Dot Notation – Lewis Dots – shows valence electrons octet rule Electron dot notation

  11. Rules for Drawing Lewis Structures • Determine the type and number of atoms • Determine the total number of valence electrons available to bond • Draw skeletal structure (least electronegative atom is usually central) • Draw bonds • Place extra electrons to fill octet • Check or recount electrons to make sure everything is “HAPPY”

  12. Simple Structures

  13. More Single Bonds

  14. Multiple Covalent Bonds

  15. Resonant Structures – two or more possible structures for the same molecule Ozone Nitrate Ion

  16. Section3 Ionic Bonding and Ionic Compounds Chapter 6 Objectives • Compare a chemical formula for a molecular compounds with one for an ionic compound. • Discuss the arrangements of ions in crystals. • Definelattice energy and explain its significance. • List and compare the distinctive properties of ionic and molecular compounds. • Write the Lewis structure for a polyatomic ion given the identity of the atoms combined and other appropriate information. (we already worked on this one)

  17. Ionic Compounds • Ionic Compounds – positive and negative ions combined in such a way that the net charge is 0 • Formula Unit – simplest collection (ratio) of atoms in an ionic compound formula unit Section 3Ionic Bonding and Ionic Compounds

  18. I o n i cB o n d s

  19. Crystal Lattice – orderly arrangement of an ionic crystal

  20. Basic Lattice Systems

  21. Covalent Molecules Low melting points Low boiling points insulators Ionic Compounds High melting points High boiling points Conductors when molten Properties of Bond Types

  22. Polyatomic Ions – a group of covalently bonded atoms that has a charge nitrate ion, NO3-

  23. Section4 Metallic Bonding Chapter 6 Objectives • Describe the electron-sea model of metallic bonding, and explain why metals are good electrical conductors. • Explain why metal surfaces are shiny. • Explain why metals are malleable and ductile but ionic-crystalline compound are not.

  24. Metallic Bonding – attraction between metals with delocalized electrons – (not a real bond) • Properties • Good conductors of heat and electricity • Good reflectors (shiny) • Malleable • Ductile • High tensile strength Section 4Metallic Bonding

  25. Section5 Molecular Geometry Chapter 6 Objectives • Explain VSEPR theory. • Predict the shapes of molecules or polyatomic ions using VSEPR theory. • Explain how the shapes of molecules are accounted for by hybridization theory.

  26. Section5 Molecular Geometry Chapter 6 Objectives, continued • Describe dipole-dipole forces, hydrogen bonding, induced dipoles, and London dispersion forces and their effects on properties such as boiling and melting points. • Explain the shapes of molecules or polyatomic ions using VSEPR theory.

  27. Molecular Geometry VSEPR Theory • – Valence Shell Electron Pair Repulsion Theory – Predicts shapes of molecules based on the repulsion of unshared pairs of electrons (lone pairs) Section 5Molecular Geometry

  28. Examples VSEPR

  29. Examples with Bond Angles

  30. Hybridization – mixing of two or more orbitals of similar energies to make orbitals with the same energy Hybridization of Carbon forms 4 equiv. orbitals

  31. Polar Molecules – molecules that have a positive and negative ends due to the arrangement of bonds and lone pairs H F

  32. Intermolecular Forces forces of attraction between molecules that are weaker than bonds attraction between polar molecules Dipole – Dipole Forces attraction that occurs between hydrogen atoms in a compound and strongly electronegative atoms with lone pairs of electrons (oxygen, nitrogen, etc.) Hydrogen Bond very small, short lived attractions and repulsions caused by the motion of electrons London Dispersion Forces

  33. Dipole -Dipole Hydrogen Bonding London Dispersion Forces

  34. Homework • Page 209-211 • Numbers 3,5,6,15,16,20,21,23,24,25,28,29,31,38,46 48,49

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