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Molecular Structures

Molecular Structures. Ch. 9. Types of Chemical Bonds. Ionic bond - the taking of electrons between atoms Covalent bond - the sharing of electrons between atoms Why do some elements take electron s and some share?. Electronegativity Difference ( Δ EN).

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Molecular Structures

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  1. Molecular Structures Ch. 9

  2. Types of Chemical Bonds • Ionic bond- the taking of electrons between atoms • Covalent bond- the sharing of electrons between atoms • Why do some elements take electrons and some share?

  3. Electronegativity Difference (ΔEN) • Electronegativity电负性- the pull an atom has on electrons while bonding • Periodic trend: • Increase moving up a group and across the periodic table from left to right • Shielding effect屏蔽效应- valence electrons feel less pull by the nucleus because lower level electrons shield them • Most electronegative element is Fluorine • ΔEN- difference in electronegative values of the elements forming the bond

  4. Calculating ΔEN • Always subtract the small EN value from the larger EN value HF ---> H= 2.1; F= 4.0 ---> 4.0-2.1= 1.9 HF= very polar covalent bond OF2---> ? Fe2O3 --->? O= 3.5; F= 4.0 ---> 4.0-3.5= 0.5 ; moderately polar covalent O= 3.5; Fe= 1.8 ---> 3.5-1.8= 1.7; very polar covalent

  5. Polar Bonds • Polar covalent bond极性共价健- a covalent bond with unequal sharing of the electrons • Creates a polar molecule (H2O) • Molecule has dipoles偶极子(δ+ and δ– ends) • Non-polar covalent bonds- a covalent bond with equal sharing of the electrons • All diatomic molecules (H2, F2, Cl2, etc…)

  6. Intermolecular Forces • Forces between molecules but not actual bonds • van der Waals forces范德华力- weak attractions between molecules • London dispersion forces伦敦色散力 • Weakest connecting force between non-polar molecules • Attraction/repulsion between nucleus and electrons • Dipole-dipole forces取向力 • Weak bonds between dipoles of polar molecules • Hydrogen bonds氢键- weak dipole force between hydrogen and electronegative elements • Strongest intermolecular force; strength in numbers

  7. Molecular Shapes • VSEPR Theory价层电子对互斥理论: • Valence Shell Electron Pair Repulsion • Valence electrons repel each other as much as possible • Using the # of lone pairs (non-bonding electrons) and # of bonds around the central atom, one can predict the 3D shape of a molecule • Double and triple bonds only count as 1 bond • Lewis-dot diagrams help organize the electrons

  8. VSEPR Models • Linear • 2 bond pairs of electrons • 0 lone pairs • Bond angle of 180o • Common compounds • CO2, CaH2, HgCl2 180o

  9. VSEPR Models • Bent (120o) • 2 bond pairs of electrons • 1 lone pairs • Bond angle of 120o • Common compounds • O3, SO2, etc… 120o

  10. VSEPR Models • Bent (105o) • 2 bond pairs of electrons • 2 lone pairs • Bond angle of 105o • Common compounds • H2O, OF2, etc… 105o

  11. VSEPR Models • TrigonalPlanar平面三角形 • 3 bond pairs of electrons • 0 lone pairs • Bond angle of 120o • Common compounds • BF3, SO3, NO3-, etc… 120o

  12. VSEPR Models • TrigonalPyramidal三角锥 • 3 bond pairs of electrons • 1 lone pairs • Bond angle of 107o • Common compounds • NH3, PCl3, etc… 107o

  13. VSEPR Models • Tetrahedral四面体的 • 4 bond pairs of electrons • 0 lone pairs • Bond angle of 109.5o • Common compounds • CH4, SO42-, etc… 109.5o

  14. VSEPR Models • TrigonalBipyramidal三角双锥 • 5 bond pairs of electrons • 0 lone pairs • Bond angle of 120 and 90o • Common compounds • PF5, PCl5, etc… 90o 120o

  15. VSEPR Models • Octahedral八面体的 • 6 bond pairs of electrons • 0 lone pairs • Bond angle of 90o • Common compounds • SF6, SeF6 etc… 90o

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