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% Document Info
\author{Katherine Doucet (katherinedoucet)}
\pdfinfo{
  /Title (chemistry-vsepr.pdf)
  /Creator (Cheatography)
  /Author (Katherine Doucet (katherinedoucet))
  /Subject (Chemistry VSEPR Cheat Sheet)
}

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\noindent
\begin{multicols}{3}
\begin{tabulary}{5.8cm}{C}
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    \vspace{-7pt}
    {\parbox{\dimexpr\textwidth-2\fboxsep\relax}{\noindent
        \hspace*{-6pt}\includegraphics[width=5.8cm]{/web/www.cheatography.com/public/images/cheatography_logo.pdf}}
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\begin{tabulary}{11cm}{L}
    \vspace{-2pt}\large{\bf{\textcolor{DarkBackground}{\textrm{Chemistry VSEPR Cheat Sheet}}}} \\
    \normalsize{by \textcolor{DarkBackground}{Katherine Doucet (katherinedoucet)} via \textcolor{DarkBackground}{\uline{cheatography.com/171479/cs/36037/}}}
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\noindent
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  \mymulticolumn{2}{p{5.377cm}}{\bf\textcolor{white}{Cheatographer}}  \\
  \vspace{-2pt}Katherine Doucet (katherinedoucet) \\
  \uline{cheatography.com/katherinedoucet} \\
  \end{tabulary}
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\begin{tabulary}{5.8cm}{L}
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  \mymulticolumn{1}{p{5.377cm}}{\bf\textcolor{white}{Cheat Sheet}}  \\
   \vspace{-2pt}Not Yet Published.\\
   Updated 12th December, 2022.\\
   Page {\thepage} of \pageref{LastPage}.
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  \mymulticolumn{1}{p{5.377cm}}{\bf\textcolor{white}{Sponsor}}  \\
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  \vspace{-5pt}
  %\includegraphics[width=48px,height=48px]{dave.jpeg}
  Measure your website readability!\\
  www.readability-score.com
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\begin{document}
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\begin{multicols*}{3}

\begin{tabularx}{5.377cm}{x{2.4885 cm} x{2.4885 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{2}{x{5.377cm}}{\bf\textcolor{white}{Molecular Geometry}}  \tn
% Row 0
\SetRowColor{LightBackground}
VSEPR & electron pairs in the valence shell of an atom repel one another and will arrange themselves to be as far apart as possible, minimizing the repulsive interactions between them \tn 
% Row Count 9 (+ 9)
% Row 1
\SetRowColor{white}
electron domain & a lone pair or a bond, regardless of whether the bond is single, double, or triple \tn 
% Row Count 14 (+ 5)
% Row 2
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electron domain geometry (EG) & arrangement of electron domains (bonds and lone pairs) around the central atom \tn 
% Row Count 18 (+ 4)
% Row 3
\SetRowColor{white}
molecular geometry (MG) & arrangement of bonded atoms \tn 
% Row Count 20 (+ 2)
% Row 4
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\mymulticolumn{2}{x{5.377cm}}{when there are no lone pairs on any of the the central atoms, MG is the same as EG.} \tn 
% Row Count 22 (+ 2)
% Row 5
\SetRowColor{white}
bond angle & angle between two adjacent bonds in a molecule or polyatomic ion \tn 
% Row Count 26 (+ 4)
% Row 6
\SetRowColor{LightBackground}
AB5 molecules contain 2 bond angles because positions occupies by bonds in a trigonal bipyramid are not all equal. & 1) equatorial: 3 bonds arranged in a trigonal plane 2) axial: 2 bonds that form an axis perpendicular to the trigonal plane \tn 
% Row Count 33 (+ 7)
\hhline{>{\arrayrulecolor{DarkBackground}}--}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{5.377cm}{x{2.4885 cm} x{2.4885 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{2}{x{5.377cm}}{\bf\textcolor{white}{Valence Bond Theory}}  \tn
% Row 0
\SetRowColor{LightBackground}
valence bond theory & bonds form between atoms when atomic orbitals overlap, thus allowing the atoms to share valence electrons. \tn 
% Row Count 6 (+ 6)
% Row 1
\SetRowColor{white}
 & each of the overlapping atomic orbitals must contain a single, unpaired electron. the two electrons shared by the bonded atoms must have opposite spins. \tn 
% Row Count 14 (+ 8)
% Row 2
\SetRowColor{LightBackground}
 & the nuclei of both atoms are attracted to the shared pair of electrons. \tn 
% Row Count 18 (+ 4)
% Row 3
\SetRowColor{white}
 & the mutual attraction for the shared electrons holds the atoms together. \tn 
% Row Count 22 (+ 4)
% Row 4
\SetRowColor{LightBackground}
valence bond theory explains why covalent bonds form. & a covalent bond will form between two atoms if the potential energy of the resulting molecule is lower than the combined potential energies of the isolated atoms. \tn 
% Row Count 31 (+ 9)
\end{tabularx}
\par\addvspace{1.3em}

\vfill
\columnbreak
\begin{tabularx}{5.377cm}{x{2.4885 cm} x{2.4885 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{2}{x{5.377cm}}{\bf\textcolor{white}{Valence Bond Theory (cont)}}  \tn
% Row 5
\SetRowColor{LightBackground}
 & formation of a covalent bond gives off energy; energy must be supplies to a molecule to break covalent bonds. \tn 
% Row Count 6 (+ 6)
\hhline{>{\arrayrulecolor{DarkBackground}}--}
\SetRowColor{LightBackground}
\mymulticolumn{2}{x{5.377cm}}{summary of valence bond theory: a bond forms when slightly occupied atomic orbitals on two atoms overlap. the two electrons shared in the region of orbital overlap must be of opposite spin. formation of a bond results in a lower potential energy for the system.}  \tn 
\hhline{>{\arrayrulecolor{DarkBackground}}--}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{5.377cm}{x{1.04425 cm} x{1.04425 cm} x{1.04425 cm} x{1.04425 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{4}{x{5.377cm}}{\bf\textcolor{white}{Intermolecular Forces}}  \tn
% Row 0
\SetRowColor{LightBackground}
\seqsplit{intermolecular} forces & \seqsplit{electrostatic} \seqsplit{attractions} between opposite charges or partial charges. & particles in the condensed phases (solids and liquids) are held together by \seqsplit{intermolecular} forces. &  \tn 
% Row Count 10 (+ 10)
% Row 1
\SetRowColor{white}
van der waals forces & \seqsplit{intermolecular} forces acting between atoms or molecules in a pure \seqsplit{substance.} & include \seqsplit{dipole-dipole} \seqsplit{interactions} (which include hydrogen bonding) and \seqsplit{dispersion} forces. &  \tn 
% Row Count 19 (+ 9)
% Row 2
\SetRowColor{LightBackground}
\seqsplit{dipole-dipole} \seqsplit{interactions} & \seqsplit{attractive} forces that act between polar \seqsplit{molecules.} & partial positive charge on one molecule is attracted to the partial negative charge on the \seqsplit{neighboring} molecule. & the larger the dipole, the larger the \seqsplit{attractive} force. \tn 
% Row Count 31 (+ 12)
\end{tabularx}
\par\addvspace{1.3em}

\vfill
\columnbreak
\begin{tabularx}{5.377cm}{x{1.04425 cm} x{1.04425 cm} x{1.04425 cm} x{1.04425 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{4}{x{5.377cm}}{\bf\textcolor{white}{Intermolecular Forces (cont)}}  \tn
% Row 3
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hydrogen bonding & a strong type of \seqsplit{dipole-dipole} \seqsplit{interaction} that occurs in molecules \seqsplit{containing} hydrogen bonded to a small, highly \seqsplit{electronegative} atom, such as nitrogen, oxygen, or fluorine. & within a series of hydrogen compounds of group 14, the boiling point increases with \seqsplit{increasing} molar mass. & for groups 15-17, the same trend is observed for all but the smallest member of each series, which has an \seqsplit{irregularly} high boiling point. \tn 
% Row Count 18 (+ 18)
% Row 4
\SetRowColor{white}
\seqsplit{dispersion} forces & \seqsplit{attractive} forces that act between all molecules, nonpolar and polar. & forces between an \seqsplit{instantaneous} dipole (fleeting, temporary dipole) and induced dipoles \seqsplit{(instantaneous} dipoles can induce \seqsplit{neighboring} dipoles). & when a nonpolar molecule acquires an \seqsplit{instantaneous} dipole, it is \seqsplit{polarized.} \tn 
% Row Count 33 (+ 15)
\end{tabularx}
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\vfill
\columnbreak
\begin{tabularx}{5.377cm}{x{1.04425 cm} x{1.04425 cm} x{1.04425 cm} x{1.04425 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{4}{x{5.377cm}}{\bf\textcolor{white}{Intermolecular Forces (cont)}}  \tn
% Row 5
\SetRowColor{LightBackground}
magnitude of \seqsplit{dispersion} forces depends on how mobile the electrons in the molecule are. & in small molecules, the electrons are \seqsplit{relatively} close to the nuclei and cannot move about very freely. thus, the electron \seqsplit{distribution} is not easily \seqsplit{polarized.} & in larger molecules, the electrons are farther away from the nucleus and can move about more freely. thus, electron \seqsplit{distribution} is easily polarized, resulting in larger \seqsplit{instantaneous} dipoles, larger induced dipoles, and larger \seqsplit{intermolecular} forces overall. & the more valence electrons a compound has, the more easily polarized it is. \tn 
% Row Count 26 (+ 26)
% Row 6
\SetRowColor{white}
if a compound is polar, \seqsplit{dipole-dipole} forces and \seqsplit{dispersion} forces are acting on it. & if a molecule is polar with a hydrogen bonded to fluorine, nitrogen, or oxygen, \seqsplit{dipole-dipole} \seqsplit{(including} hydrogen) and \seqsplit{dispersion} forces are acting on it. & if the compound is nonpolar, only \seqsplit{dispersion} forces are acting on it. &  \tn 
% Row Count 42 (+ 16)
\end{tabularx}
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\vfill
\columnbreak
\begin{tabularx}{5.377cm}{x{1.04425 cm} x{1.04425 cm} x{1.04425 cm} x{1.04425 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{4}{x{5.377cm}}{\bf\textcolor{white}{Intermolecular Forces (cont)}}  \tn
% Row 7
\SetRowColor{LightBackground}
\seqsplit{ion-dipole} \seqsplit{interactions} & \seqsplit{attraction} between ions and polar molecules in \seqsplit{solutions.} & magnitude depends on the charge and size of the ion and on the dipole moment and size of the polar molecule. & cations interact more strongly with dipoles than anions because they tend to be smaller. an ion with higher charge and smaller size will interact more strongly with water \seqsplit{molecules.} \tn 
% Row Count 19 (+ 19)
% Row 8
\SetRowColor{white}
strongest force to weakest force: & \seqsplit{ion-dipole}, hydrogen, \seqsplit{dipole-dipole}, \seqsplit{dispersion} &  &  \tn 
% Row Count 24 (+ 5)
\hhline{>{\arrayrulecolor{DarkBackground}}----}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{5.377cm}{X}
\SetRowColor{DarkBackground}
\mymulticolumn{1}{x{5.377cm}}{\bf\textcolor{white}{Hybridization of Multiple Bonds}}  \tn
% Row 0
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\mymulticolumn{1}{x{5.377cm}}{sigma bonds} \tn 
% Row Count 1 (+ 1)
% Row 1
\SetRowColor{white}
\mymulticolumn{1}{x{5.377cm}}{pi bonds} \tn 
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% Row 2
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\mymulticolumn{1}{x{5.377cm}}{singly occupied p orbitals give rise to multiple bonds.} \tn 
% Row Count 4 (+ 2)
% Row 3
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\mymulticolumn{1}{x{5.377cm}}{pi bonds restrict the rotation of a molecule in a way that sigma bonds do not.} \tn 
% Row Count 6 (+ 2)
\hhline{>{\arrayrulecolor{DarkBackground}}-}
\end{tabularx}
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\begin{tabularx}{5.377cm}{x{2.4885 cm} x{2.4885 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{2}{x{5.377cm}}{\bf\textcolor{white}{Deviation from Ideal Bond Angles}}  \tn
% Row 0
\SetRowColor{LightBackground}
lone pairs take up more space than bonding pairs. & a lone pair on the central atom is attracted only to the nucleus of that atom. a bonding pair is simultaneously attracted to the nuclei of both the bonding atoms. lone pairs have more freedom to spread out and greater capacity to repel other electron domains. \tn 
% Row Count 13 (+ 13)
% Row 1
\SetRowColor{white}
multiple bonds repel more strongly than single bonds because they contain more electron density. & multiple bonds are shorter than single bonds. \tn 
% Row Count 18 (+ 5)
\hhline{>{\arrayrulecolor{DarkBackground}}--}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{5.377cm}{x{2.4885 cm} x{2.4885 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{2}{x{5.377cm}}{\bf\textcolor{white}{Bonding Theory}}  \tn
% Row 0
\SetRowColor{LightBackground}
in species that can be represented by two or more resonance structures, the pi bonds are delocalized, meaning that they are spread out over the molecule and not constrained to just two atoms. & localized bonds are those constrained to two atoms. \tn 
% Row Count 10 (+ 10)
\hhline{>{\arrayrulecolor{DarkBackground}}--}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{5.377cm}{x{1.8308 cm} x{2.2885 cm} p{0.4577 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{3}{x{5.377cm}}{\bf\textcolor{white}{Hybridization of Atomic Orbitals}}  \tn
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hybridization & atomic orbitals mix to form hybrid orbitals. &  \tn 
% Row Count 3 (+ 3)
% Row 1
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hybrid orbitals & orbitals formed by hybridization of some combination of s, p, or d atomic orbitals. &  \tn 
% Row Count 8 (+ 5)
% Row 2
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each sp hybrid orbital has one small lobe and one large lobe. & they are oriented in opposite directions with a 180 degree angle between them. &  \tn 
% Row Count 12 (+ 4)
% Row 3
\SetRowColor{white}
\mymulticolumn{3}{x{5.377cm}}{number of electron domains is equal to the number of hybrid orbitals which is also equal to the number of atomic orbitals needed to mix.} \tn 
% Row Count 15 (+ 3)
% Row 4
\SetRowColor{LightBackground}
elements in the third period of the periodic table and beyond do not obey the octet rule because they have d orbitals that can hold additional electrons. & in molecules where there are more than 4 electron domains on the central atom, include d orbitals in hybridization. &  \tn 
% Row Count 25 (+ 10)
% Row 5
\SetRowColor{white}
 & sp\textasciicircum{}3\textasciicircum{}d have shapes similar to sp, sp\textasciicircum{}2\textasciicircum{}, and sp\textasciicircum{}3\textasciicircum{} orbitals - one large lobe and one small lobe. &  \tn 
% Row Count 30 (+ 5)
\hhline{>{\arrayrulecolor{DarkBackground}}---}
\SetRowColor{LightBackground}
\mymulticolumn{3}{x{5.377cm}}{do not use hybrid orbitals to predict molecular geometries but to explain geometries that are already known.}  \tn 
\hhline{>{\arrayrulecolor{DarkBackground}}---}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{5.377cm}{x{2.4885 cm} x{2.4885 cm} }
\SetRowColor{DarkBackground}
\mymulticolumn{2}{x{5.377cm}}{\bf\textcolor{white}{Molecular Geometry and Polarity}}  \tn
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\SetRowColor{LightBackground}
polar molecule & a bond between two atoms of different electronegativities \tn 
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% Row 1
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a molecule bonded with a highly electronegative element will be polar. & highly electronegative elements: fluorine, nitrogen, oxygen, or chlorine \tn 
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\mymulticolumn{2}{x{5.377cm}}{if the central atom has lone pairs, molecule is most likely polar} \tn 
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if vectors cancel each other, molecule is nonpolar. & if vectors do not cancel each other (are not equal) the molecule is polar. \tn 
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structural isomers & molecules that have the same chemical formula but have different arrangement of atoms \tn 
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\hhline{>{\arrayrulecolor{DarkBackground}}--}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{5.377cm}{X}
\SetRowColor{DarkBackground}
\mymulticolumn{1}{x{5.377cm}}{\bf\textcolor{white}{Molecular Orbital Theory}}  \tn
% Row 0
\SetRowColor{LightBackground}
\mymulticolumn{1}{x{5.377cm}}{molecular orbital theory} \tn 
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% Row 1
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\mymulticolumn{1}{x{5.377cm}}{molecular orbitals} \tn 
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\mymulticolumn{1}{x{5.377cm}}{} \tn 
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\SetRowColor{white}
\mymulticolumn{1}{x{5.377cm}}{} \tn 
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% Row 4
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\mymulticolumn{1}{x{5.377cm}}{} \tn 
% Row Count 2 (+ 0)
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\mymulticolumn{1}{x{5.377cm}}{bonding molecular orbital} \tn 
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\mymulticolumn{1}{x{5.377cm}}{antibonding molecular orbital} \tn 
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\mymulticolumn{1}{x{5.377cm}}{sigma molecular orbitals} \tn 
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\mymulticolumn{1}{x{5.377cm}}{bond order} \tn 
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\mymulticolumn{1}{x{5.377cm}}{pi molecular orbitals} \tn 
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% Row 10
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\mymulticolumn{1}{x{5.377cm}}{molecular orbitals resulting from the combination of p atomic orbitals are higher in energy than the molecular orbitals resulting from the combination of s atomic orbitals.} \tn 
% Row Count 11 (+ 4)
\hhline{>{\arrayrulecolor{DarkBackground}}-}
\end{tabularx}
\par\addvspace{1.3em}

\begin{tabularx}{5.377cm}{X}
\SetRowColor{DarkBackground}
\mymulticolumn{1}{x{5.377cm}}{\bf\textcolor{white}{Hybridization Chart}}  \tn
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\mymulticolumn{1}{x{5.377cm}}{Number of Electron Domains on Central Atom} \tn 
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\mymulticolumn{1}{x{5.377cm}}{hybrid orbitals} \tn 
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\mymulticolumn{1}{x{5.377cm}}{geometry} \tn 
% Row Count 3 (+ 1)
\hhline{>{\arrayrulecolor{DarkBackground}}-}
\end{tabularx}
\par\addvspace{1.3em}


% That's all folks
\end{multicols*}

\end{document}