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IB DP1 Chemistry HL Bonding

IB DP1 Chemistry HL Bonding. What makes atoms join together to make compounds?. Topic 14 : Bonding. 14.1 Shapes of molecules and ions 1 hour 14.1.1 Predict the shape and bond angles for species with five and six negative charge centres using the VSEPR theory. 14.2 Hybridization

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IB DP1 Chemistry HL Bonding

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  1. IB DP1 ChemistryHL Bonding What makes atoms join together to make compounds?

  2. Topic 14: Bonding 14.1 Shapes of molecules and ions 1 hour 14.1.1 Predict the shape and bond angles for species with five and six negative charge centres using the VSEPR theory. 14.2 Hybridization 2 hours 14.2.1 Describe σ and π bonds. 14.2.1 Describe σ and π bonds. 14.2.2 Explain hybridization in terms of the mixing of atomic orbitals to form new orbitals for bonding. 14.2.3 Identify and explain the relationships between Lewis structures, molecular shapes and types of hybridization (sp, sp2 and sp3). 14.3 Delocalization of electrons 2 hours 14.3.1 Describe the delocalization of π electrons and explain how this can account for the structures of some species.

  3. Schrodinger wave equation

  4. Whichenergylevel is an electron in? 2p1 2nd quantum number 1st quantum number 3rd quantum number

  5. Electronorbitalshapes (2nd quantum number) http://chemwiki.ucdavis.edu/Physical_Chemistry/Quantum_Mechanics/Atomic_Theory/Electrons_in_Atoms/Electronic_Orbitals

  6. Electronenergylevels

  7. Hybridization atoms  circular electron shells 2,8,8,…orbitals s,p,d,f,… • s, p, d, f orbitals only for single atoms in gaseous state  hybridization • electron orbitals change shape (and energy) during bonding

  8. σ-bond • strongest form of covalent bond • orbitals overlap on line between nuclei • commonly s+s, pz+pz, s+pz Image: http://en.wikipedia.org/wiki/Pi_bond

  9. π-bond • orbital overlap not on line between nuclei • usually weaker than sigma bonds • stop rotation Image: http://en.wikipedia.org/wiki/Pi_bond

  10. Orbital shapes of spdf orbitals and hybrid ized orbitals Image: http://en.wikipedia.org/wiki/Pi_bond

  11. Electronic configuration of carbon 1s22s22p2

  12. Methane sp3 hybridization • 2s and 2p3 orbitals hybridize

  13. Ethane, ethene and ethyne

  14. Carbon-carbon bonds • Describe and explain the change in bond energy • Describe and explain the change in bond length

  15. Single bond (ethane) • one axial C-C s -Bond • Hybridisation: ones-orbital and three p-orbitals  four sp3-orbitals • The sp3-orbitals have a tetrahedral shape (109.5o).

  16. Double bond(ethene) • one axials -bond and one offset p-bond • Hybridisation: one s-orbital and two p-orbitals threesp2-orbitals • The sp2-orbitals have a trigonal planar shape, 120o

  17. Triple bond • One axial s -bond and two offset p -bonds • Hybridisation: One s-orbital and one p-orbital Two sp-orbitals • The sp-orbitals give a linear shape

  18. The shape of the hybrids corresponds to the structure given by VSEPR / Lewis structure. • Ethane : Ethene : Ethyne sp3: sp2: sp Ammonia: sp3 Water: sp3

  19. Electrons not associatedwith a particular atom or bondaredelocalized metallic bond

  20. Benzene, C6H6 ring B A • The p-bond in the double bond can switch place Electrons are delocalized. • A and B: resonance structures. • C: resonance hybrid. Molecule gains resonance energy by delocalizing electrons. C Image: http://commons.wikimedia.org/wiki/File:Benzene_resonance_structures.png

  21. Resonance bond • Resonance describes delocalized electrons within some molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis formula. • A molecule or ion with such delocalized electrons is represented by several resonance structures. Draw the structures of NO3-, CO32-, O3

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