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Formed when two elements share valence electrons so that both gain full valence shells.
Bonds can be single (two shared electrons)
double (four shared electrons)
triple (six shared electrons). |
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| Electrons in the outermost occupied electron shell of an atom. |
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An atom is most stable when its valence shell is completely full; an atom can obtain a full valence shell by bonding with other atoms.
Does not always apply to atoms with electrons in d orbitals. |
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Formed when the electro negativities of the bonded atoms are significantly different.
Both shared electrons are almost completely associated with the more electronegative atom, which becomes an anion (negatively charged) while the other atom becomes a cation (positively charged). |
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| Half the number of electrons shared in a covalent bond. The higher the bond order, the stronger and shorter the bond compared to a lower order bond between the same two atoms; e.g., a C=C bond (bond order of 2) is stronger and shorter than a C–C bond (bond order of 1). |
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| In bonds between atoms of unequal electro negativities, shared electrons are more likely to be found near the more electronegative atom. Such a bond is said to be polar. |
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| Ability of an atom in a molecule to attract electrons. Increases traveling across a period from left to right in the Periodic Table of the Elements, and decreases traveling down a group in the Periodic Table. |
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| Energy required to break a chemical bond. The higher the energy, the stronger the bond. |
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| Valence shell electron pair repulsion (VSEPR) theory |
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| Model for predicting molecular geometry. Based on the idea that, in a molecule, electron pairs arrange themselves as far apart as possible to minimize electron-electron repulsion. |
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| Theory of bonding that states that covalent bonds form through the spatial overlap of orbitals containing valence electrons. Valence bond theory is consistent with the geometric predictions of VSEPR theory. |
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| Formation of hybrid orbitals, which are mixtures of individual atomic orbitals. |
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| Bonds formed by the head-on overlap of sp, sp2, or sp3 hybrid orbitals with each other or with hydrogen 1s orbitals. |
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| Bonds formed by the sideways overlap of p orbitals |
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| Electrons in a (π) orbital (either bonding or nonbonding). |
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| A molecular representation that depicts covalent bonds and nonbonding valence electrons. |
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| Some Lewis structures have multiple, distinct, but equivalent ways of arranging multiple bonds and electrons, while still obeying the Octet Rule and maintaining the connectivity of atoms in the molecule. Each equivalent arrangement is called this term. |
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| The formal charge on an atom equals the number of valence electrons in the unbonded atom minus both the number of lone-pair electrons the atom has and the number of covalent bonds to the atom. |
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| Shorthand method of representing organic molecules in which some or all bonds are not explicitly drawn. |
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Representation of an organic molecule in which:
carbon atoms and hydrogen atoms bonded to carbon are not explicitly drawn.
carbon-carbon bonds are represented by lines, with carbon atoms assumed to be present at the intersection of any two lines and at the ends of lines.
any carbon not explicitly making four covalent bonds is implicitly bonded to however many hydrogen atoms are necessary to give that carbon four bonds. |
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| Shows the three-dimensional arrangement of atoms in a molecule. Dotted wedges indicate bonds pointing into the plane of the paper; solid wedges indicate bonds rising from the plane of the paper; and solid lines indicate bonds that are in the plane of the paper. |
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