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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 - PowerPoint PPT Presentation
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IB DP1 ChemistryHL Bonding
What makes atoms join together to make compounds?
Topic 14: Bonding14.1 Shapes of molecules and ions1 hour14.1.1 Predict the shape and bond angles for species with five and six negative charge centres using the VSEPR theory.14.2 Hybridization2 hours14.2.1 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 electrons2 hours14.3.1 Describe the delocalization of π electrons and explain how this can account for the structures of some species.
Schrodinger wave equation
Which energy level is an electron in?
2p11st quantum number
2nd quantum number
3rd quantum number
Electron orbital shapes (2nd quantum number)
http://chemwiki.ucdavis.edu/Physical_Chemistry/Quantum_Mechanics/Atomic_Theory/Electrons_in_Atoms/Electronic_Orbitals
Electron energy levels
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
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
π-bond
orbital overlap not on line between nuclei usually weaker than sigma bonds stop rotation
Image: http://en.wikipedia.org/wiki/Pi_bond
Orbital shapes of spdf orbitals and hybrid ized orbitals
Image: http://en.wikipedia.org/wiki/Pi_bond
Electronic configuration of carbon
1s2 2s2 2p2
Methane sp3 hybridization 2s and 2p3 orbitals hybridize
Ethane, ethene and ethyne
Carbon-carbon bonds Bond
typeBond energy (kJ/mol)
Bond length (pm)
Hybrid orbitals
ethane single 348 154
ethene double 612 134
ethyne triple 837 120
• Describe and explain the change in bond energy• Describe and explain the change in bond length
Single bond (ethane) one axial C-C s -Bond Hybridisation: one s-orbital and three p-orbitals four
sp3-orbitals The sp3-orbitals have a tetrahedral shape (109.5o).
Double bond (ethene) one axial s -bond and one offset p-bond Hybridisation: one s-orbital and two p-orbitals three
sp2-orbitals The sp2-orbitals have a trigonal planar shape, 120o
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
The shape of the hybrids corresponds to the structure given by VSEPR / Lewis structure.
Ethane : Ethene : Ethyne sp3 : sp2 : sp
Ammonia: sp3
Water: sp3
Electrons not associated with a particular atom or bond are delocalized
metallic bond
Benzene, C6H6 ring
Image: http://commons.wikimedia.org/wiki/File:Benzene_resonance_structures.png
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.
A
C
B
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-, CO3
2-, O3