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14 Cards in this Set
- Front
- Back
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Pure Covalent Bond |
• e- shared equally between 2 identical atoms • diatomic molecules |
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Polar Covalent Bond |
• e- shared unequally between 2 atoms • The more electro- atom will hold the e- more tightly |
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Lewis Structures |
• Most important aspect is often satisfying the octet rule • Although there will be a few rare exceptions when forced |
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Molecular Geometry |
• The shape of only bonded atoms • Most consider the electron domain as this will determine the bond angles & shape |
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VSEPR Valence She’ll Electron Pair Repulsive |
• The shapes molecules form are to separate the valence electrons as far as possible • The number of groups that contain electrons will determine the shape of the molecule |
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No polar Covalent |
e- shared close to same from both atoms, so no partial charge forms (CH,CS) |
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Multiple Bonding |
There are times where multiple bonds will form between two atoms in order to complete octet of both atoms |
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Double Bonds |
• 2 pairs shared e- • Shorter & stronger than single bonds Ex: CO2 |
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Bond Angles |
- Some have two angles • This is due to lone pair e- taking up more space than bonded atoms • So the angles between the atoms become smaller by 2 to 4° • On test indicate by stating slightly less than the normal angle |
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Covalent Bonds |
• Bond by sharing e- • Not full e- transfer like ionic bonds • Often between 2 non-metals • Most atoms can form multiple covalent bonds |
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Lewis Structures |
• Simple models of covalent bonding structures • Deals with valence e- only |
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Triple Bonds |
• 3 pairs shared e- • Shorter & stronger than double bonds Ex: N2 |
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Electron Domain |
• Bonded atoms + lone pairs e- • Need to be counted both for the correct structure |
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Polar Covalent Bond |
• This creates a diplo partial (+) and (-) charges to the bond * More electronegative will have 8- charge |
