1 / 25

Ch. 6—Chemical Bonding

Ch. 6—Chemical Bonding. 6-1 Introduction to Chemical Bonding 6-2 Covalent Bonding and Molecular Compounds 6-3 Ionic Bonding and Ionic Compounds 6-4 Metallic Bonding 6-5 Molecular Geometry. 6-1 Chemical Bonding.

denis
Download Presentation

Ch. 6—Chemical Bonding

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. Ch. 6—Chemical Bonding 6-1 Introduction to Chemical Bonding 6-2 Covalent Bonding and Molecular Compounds 6-3 Ionic Bonding and Ionic Compounds 6-4 Metallic Bonding 6-5 Molecular Geometry

  2. 6-1 Chemical Bonding • Chemical bond– a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together • Ionic bonding– results from the electrical attraction between large numbers of cations and anions • Covalent bonding– results from the sharing of electron pairs

  3. Ionic or Covalent…or somewhere in between? Look at sample problem 6-1 on p. 163

  4. Covalent Bonds • Nonpolar-covalent– the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge (ex. H2) • Polar-covalent– the bonded atoms have an unequal attraction for the shared electrons (ex. HCl) • Polar covalent bonds have a positive and negative dipole (charge) indicated by δ+ or δ-

  5. 6-2 Molecules • Molecule– neutral group of atoms that are held together by covalent bonds • Other vocab—chemical formula, molecular compound, molecular formula, diatomic molecule

  6. Covalent Bonds • Often atoms are more stable as compounds rather than elements—when atoms get close together, there is an attraction between positive part of one atom and negative part of the other atom until they get too close and begin to repel each other • The distance that results in the smallest potential energy is the stable bond with a given bond length • Bond energy– energy required to break a chemical bond and form neutral isolated atoms

  7. Octet Rule • Hydrogen and helium only require 2 electrons to fill the first energy level • After that 8 electrons are required to have a similar electron configuration to the stable noble gases • Octet rule– chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level • Exceptions include elements like boron with 3 valence e- that tends to form bonds surrounded by six electrons or more than 8 valence e- with highly electronegative elements such as fluorine, oxygen, and chlorine

  8. Lewis Structure • Electron dot notation– an e- configuration notation in which only the valence e- of an atom of a particular element are shown, indicated by dots placed around the element’s symbol (as seen in ch. 3) • Lewis structures– formulas in which atomic symbols represent nuclei and inner-shell e-, dot-pairs or dashes between two atomic symbols represent e- pairs in covalent bonds and dots adjacent to only one atomic symbol represent unshared e- Unshared pair or lone pair H : H Look at sample problem 6-3 on p. 171

  9. Multiple Bonds • Single bond • Double bond • Triple bond See sample problem 6-4 on p. 174

  10. Resonance • Resonance– refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure

  11. 6-3 Ionic Bonds • Formula unit– the simplest collection of atoms from which an ionic compound’s formula can be established • Ionic compound– composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal • Ex. H2O-- molecule H2 diatomic element • NaCl– formula unit Na+cation; Cl- anion

  12. Formation of Ionic Compounds • NaCl– sodium is group IA with 1 valence e- and chlorine is group VIIA with 7 valence e- Therefore, Na+ and Cl- form NaCl with sodium giving it’s valence e- to chlorine to complete the octet rule The cation name is the same (sodium); the anion name ends in –ide (chloride) for the formula unit sodium chloride • Al2O3– aluminum has 3 valence e- and oxygen has 6 valence e-; to complete the octet rule, aluminum gives up all of it’s valence e- but oxygen only needs 2 e- so the extra e- goes to another oxygen atom—so that aluminum and oxygen are stable, it requires 2 aluminum atoms combined with 3 oxygen atoms to complete the octet rule

  13. Ionic vs. Molecular Compounds

  14. Polyatomic Ions • Polyatomic ion– charged group of covalently bonded atoms • Common ones to know • NH4+ ammonium • NO3- nitrate • SO42- sulfate • PO43- phosphate • Ex. NH4Cl ammonium chloride and (NH4)2O ammonium oxide are ionic compounds • Ex. K2SO4 potassium sulfate is an ionic compound

  15. 6-4 Metallic Bonding • Metallic bonding– chemical bonding that results from the attraction between metal atoms and the surrounding sea of e- • Metals tend to have vacant orbitals so the attraction of e- from other metal atoms toward the nucleus of metal atoms forms a sea of e- where when the atoms are close together the orbitals overlap and an e- can be part of different nuclei at any given time, or are delocalized • This allows metals to slide making them malleable (ability to be pounded to a shape) and ductile (drawn into a wire) • Because of the freedom of motion, metals have high electrical and thermal conductivity; because of many orbitals separated by small energy differences, metals can absorb different quanta giving metals their high luster • Metallic bond strength indicated by heat of vaporization (kj/mol)

  16. 6-5 Molecular Geometry • Molecular polarity– the uneven distribution of molecular charge • Chemical formulas only describe the number and type of atoms-- • VSEPR (valence-shell, electron-pair repulsion) theory describes molecular bond angles • Hybridization describes the orbitals that contain the valence e-

  17. VSEPR • VSEPR theory– states that repulsion between the sets of valence-level e- surrounding an atom causes these sets to be oriented as far apart as possible

  18. Predicting the Molecular Geometry • write the Lewis structure • fit the type of molecule based on basic structure • Ex. AB2 is linear (BeF2) • Ex. AB2E where E is an unshared pair is bent (SnCl2) • Ex. AB3E is trigonal-pyramidal (NH3) • Ex. AB2E2 is bent (H2O)

  19. 8.3 Hybrid Orbitals • Orbital hybridization provides information about both molecular bonding and molecular shape. • In hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals.

  20. 8.3 Hybridization Involving Single Bonds Methane, CH4, sp3

  21. 8.3 Hybrid OrbitaHybridization Involving Double Bonds Ethene, C2H4, sp2

  22. 8.3 Hybridization Involving Triple Bonds Ethyne, C2H2, sp

  23. Intermolecular Forces • Intermolecular forces—forces of attraction between molecules • Weaker than bonds that join atoms in molecules, ionic compounds, and metal atoms) • Dipole-dipole forces– the positive dipole of one atom is attracted to the negative dipole of another atom • Hydrogen bonding– hydrogen is attracted to oxygen, nitrogen, and fluorine of other molecules (hydrogen bonding very important in properties of water) • London dispersion forces– intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles (temporary dipoles created by e- continuous movement)

  24. Chapter 6 Objectives • What is a chemical bond? Why do atoms bond? • What is ionic bonding? Which two types of elements are involved in ionic bonding? • What is covalent bonding? What one type of element is involved in covalent bonding? • How can we determine if the predominant character of a bond is going to be ionic or covalent? • Predict the type of bond that could form between two atoms • What is the difference between a chemical formula and a molecular formula? • What are diatomic molecules (HOFBrNCLI) • Describe how a covalent bond is formed? • What is the relationship between bond length and bond energy? • What is the octet rule? Name some exceptions • Draw Lewis dot structures for compounds; draw resonance structures where appropriate • Briefly describe how ionic compounds are formed • Compare and contrast ionic and covalent bonding • What is metallic bonding? • How does metallic bonding explain certain unique properties of metals such as luster, ability to conduct heat and electricity, malleability and ductility • 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 • Describe various intermolecular forces of attraction (dipole-dipole, hydrogen, London dispersion forces) • Predict the type of intermolecular force that can exist between different covalent compounds

More Related