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Chemical Bonding

Chemical Bonding. Advanced Chemistry Ms. Grobsky. What is Bonding? Why do Atoms Bond?. Bonding is the interplay between interactions between atoms Energetically favored Electrons on one atom interacting with protons of another atom Energetically unfavorable

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Chemical Bonding

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  1. Chemical Bonding Advanced Chemistry Ms. Grobsky

  2. What is Bonding? Why do Atoms Bond? • Bonding is the interplay between interactions between atoms • Energetically favored • Electrons on one atom interacting with protons of another atom • Energetically unfavorable • Electrons on one atom interacting with electrons of another atom • Protons on one atom interacting with protons of another atom • A bond will form if the system can LOWER its total energy in the process

  3. Ionic Bonds • Bond between a metal cation and non-metal anion • Formula determined by ionic charges • Electron(s) transferred from cation to anion • Electrostatic in nature

  4. Ionic Bonds (Continued) • Ionic compounds form huge, repeating 3-D crystalline lattices • Ions and electrons are located at fixed positions • Strong interactions between ions • Large melting points • Solids at room temperature

  5. Covalent Bonds • Bond between two non-metals atoms • Valence electrons are shared between nuclei of bonding atoms • When shared equally, bond is called non-polar covalent • When shared unequally, bond is called polar covalent and dipoles are established • Sharing based on electronegativity of each atom in bond • Bonds can be single, double, or triple as shown by Lewis structures • Physical properties vary wildly

  6. How Do Covalent Bonds Form? • Sharing of valence electrons • Electrons in the highest occupied energy shell of the atom • TOTAL highest energy s and p electrons • Focus on ns, np, and d electrons of transition elements

  7. Single and Multiple Bonds • Single bond • One pair of electrons shared • Double bond • Two pairs of electrons shared • Triple bond • Three pairs of electrons shared

  8. Multiple Bonds and Bond Lengths • Multiple bonds increase electron density between two nuclei • Decreases nuclear repulsions while enhancing the nucleus to electron density attractions • Nuclei move closer together • Bond lengths from shortest to longest are as follows: Triple bond < Double bond < Single bond • The shorter the bond implies that atoms are held together more tightly when there are multiple bonds • Multiple bonds are stronger than single bonds

  9. How Do We Describe the Structure of Covalent Bonds? • Called the Localized Electron Model • Used to describe covalent bonds • Assumes that electrons are localized (restricted to certain areas) on an atom or the space between atoms • Lone pair electrons • Bonding pair electrons • You will learn about 2 parts of the model: • Lewis Dot structure describe valence electron arrangement • Geometry is predicted with VSEPR

  10. Lewis Dot Structures • Lewis Dot structures are also known as electron dot diagrams • These diagrams show only the valence (bonding) electrons • Unpaired (single) electrons will participate in bonding • Paired electrons will not participate in bonding • Octet Rule • Most elements obey octet rule • Each atom in a covalent bond has a TOTAL of 8 valence electrons around it • Most important requirement for the formation of a stable compound is that atoms achieve a noble gas configuration (octet) • There are EXCEPTIONS to this rule! • H – 2 electrons total • Be – 4 electrons total • B – 6 electrons total • n = 3 and above – expanded octets from d orbitals • NO, NO2, and ClO2 contain an odd number of valence electrons and thus, cannot obey octet rule

  11. Steps to Draw Lewis Dot Diagrams for Elements • Determine total number of valence electrons • Predict # of bonds by counting the number of unpaired electrons in Lewis structure

  12. Steps to Draw Lewis Dot Structures for Compounds • Determine total number of valence electrons • Add them up for BOTH compounds! • Add for anions, subtract for cations • Predict # of bonds by counting the number of unpaired electrons in Lewis structure • Least electronegative atom is the center atom • Remember the trend! • Draw a single bond , -, (2 electrons) to each atom • Subtract from total • Add lone pair electrons, :, to terminal atoms to satisfy octet rule • Extras go to central atom • If central atom is not octetand extra electrons are left unpaired, form multiple bonds! • Carbon bonded to N, O, P, S tend to form double bonds • Hydrogen is ALWAYS a terminal atom • Only makes 1 bond

  13. Ionic Compounds and Lewis Dot Structures • Ionic Lewis Dot structures are drawn exactly the same way as covalent compounds • ONE EXCEPTION – Ionic compounds only form SINGLE bonds! • Metal donates all valence electrons to non-metal

  14. Expanded Octets and Lewis Dot Structures • Sometimes, an atom is unable to form a stable compound by following the octet rule • Some atoms can make compounds using paired electrons in their inner shell (d and f-orbitals) • This causes expanded octets • Create more bonds than expected • Example: BrF3 and PCl5

  15. Polyatomic Ions and Lewis Dot Structures • Some covalently bonded atoms can have a few extra or fewer electrons, resulting in an overall charge • Negative charge (anions) – additional electrons must be added • Positive charge (cations) – electrons need to be reduced (subtract) • Examples: NH4+ and SO42-

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