Pre-university Advanced Level Chemistry: VSEPR theory - predicting molecule shapes
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Doc Brown's Advanced A Level Chemistry Theoretical Physical Chemistry Revision Notes The Shapes of Molecules and Ions and bond angles related to their Electronic Structure - mainly inorganic molecules on this page Part 1 from diatomic molecules to polyatomic molecules email doc brown - comments - query? All by structure and chemical bonding revision notes All my advanced A level inorganic chemistry notes
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If you think I've missed a molecule/ion shape that I should have covered for pre-university level, email it to me! Sub-index for 'shapes of molecules' pages
Introduction to VSEPR theory You could be given familiar and unfamiliar examples of species and asked to deduce the shape according to the valence shell electron pair repulsion (acronym VSEPR) theory principles.
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Introduction
continued
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electron pair repulsion theory and bond angle These relative electron pair repulsions have a profound effect on bond angles and molecule shape, both of which you need to be able to deduce. For most pre-university courses you should be able to explain the shapes of, and bond angles in, of simple molecules and ions with two bonding pairs of electrons OR three to six electron pairs (including both bonding and lone pairs of electrons) surrounding the central atom. The shape of a molecule is determined by the number of groups of electrons around the central atom. The 'groups' might be a non-bonding single electron, a non-bonding or bonding pair of electrons, a double pair of bonding electrons or triple pair of bonding electrons etc. The electron 'groupings' repel to minimise the potential energy of the system i.e. to make the A-B-C angle as wide as possible. The dot and cross diagrams (ox) are presented in 'Lewis style' In the diagrams the central atom is denoted by X and attached surrounding bonded atoms by Q. The bond angle is therefore based on angle between the atoms Q-X-Q. The phrase lone pair, usually refers to a pair of non-bonding electrons in the outer valence shell of the central atom of the molecule.
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Diatomic molecules These are not considered to have a 'shape' or a 'bond angle' in the context of this page, but they are useful dot and cross diagram revision based on the outer valence electrons and help you to construct Lewis dot and cross diagrams for molecules with >2 atoms. |
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H-H |
e.g. hydrogen H2
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H-Cl |
e.g. hydrogen chloride HCl, HX in
general where X = halogen
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Cl-Cl |
e.g. chlorine Cl2, iodine(I) chloride
ICl (iodine monochloride)
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O=O |
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Two groups of electrons around the central atom two bonding pairs of electrons (single bonds) or two double bond pairs give a linear shape and bond angle of 180o In these examples the electron pair geometry is the same as the molecular geometry. |
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![]() H-Be-H linear shape |
gaseous beryllium hydride BeH2 (Q = H, X = Be) Be is a group 2 element with two outer electrons and so can form two single covalent bonds - 2 bond pairs with hydrogen atoms - but no 'octet' of valence electrons on the central atom. linear shaped molecule (VSEPR theory argument) |
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![]() Cl-Be-Cl linear shape |
gaseous beryllium halides BeCl2 (X = Be, Q = F, Br, Cl or I) Be is a group 2 element with two outer electrons and so can form two single covalent bonds - 2 bond pairs with halogen atoms - but no 'octet' of valence electrons on the central atom. Similarly the VSEPR theory argument gives a linear shaped molecule for gaseous beryllium fluoride, beryllium chloride, beryllium bromide and gaseous beryllium iodide too I presume? |
valence bond dot and cross diagrams
the Lewis dot and cross electronic diagram used to predict the shape of carbon dioxide |
O=C=O linear shape molecule |
Here you have two groups of bonding electrons either side of a central atom in the carbon dioxide triatomic molecule producing a linear shape. Double bond pairs count just the same as single bond pairs in determining the shape of a molecule. |
[H3N-Ag-NH3]+ | [Ag(NH3)2]+ |
transition metal complex of co-ordination number 2: e.g. the diamminesilver(I) ion, [Ag(NH3)2]+, where the :NH3 ammonia molecule acts as an electron pair donor to form the dative covalent bond. Its not a true linear shaped complex ion because of the pyramidal ammonia ligands but the N-Ag-N bonds are linear i.e. bond angle 180o. |
H-C≡N |
HCN linear shape |
methanenitrile (hydrogen cyanide) is also a linear molecule. In this case the triple bond counts just the same as a single bond in terms of determining the shape. |
Three groups of electrons around the central atom |
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trigonal planar electron pair geometry |
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electrons: two bond pairs, one lone pair shape of molecule V or BENT, bond angle approximately 120o Does anyone know of any example? Q would have to be H. |
![]() the Lewis dot and cross electronic diagram used to predict the shape - trigonal planar electron pair geometry |
![]() V or bent shape molecule |
electrons: two bond pairs, one lone pair molecule shape V or BENT, bond angle approximately 120o Does anyone know of an example? For X=Q double bonds see sulfur dioxide SO2, NO2- and O3 involve 2 groups of bonding electrons and lone pair and VSEPR correctly predicts their V or bent shape and a Q-X-Q bond angle of ~120o. Electron pair geometry different from the molecular geometry. |
![]() the Lewis dot and cross electronic diagram used to predict the shape of boron hydride - trigonal planar electron pair geometry |
![]() trigonal planar shaped molecule |
electrons: 3 bond pairs, no lone pairs Electron pair geometry same as the molecular geometry. (molecular compounds/ions involving hydrogen)
e.g. X = B and Q = H for gaseous boron hydride BH3. Boron in Group 3/13 has three outer valence electrons each of which pairs up with a hydrogen electron - dot and cross diagram on the right. Boron is in group 3/13 with three outer valence electrons and can therefore form three single covalent bonds with hydrogen, but no 'octet' of valence electrons on the central atom.
DO NOT assume examination questions will always be based on neutral molecules! |
![]() the Lewis dot and cross electronic diagram used to predict the shape of some group III trihalides - trigonal planar electron pair geometry |
![]() trigonal planar shaped molecule/ion |
electrons: 3 bond pairs, no lone pairs Electron pair geometry same as the molecular geometry. (molecular compounds/ions NOT involving hydrogen) The VSEPR theory argument gives the molecule shape TRIGONAL PLANAR: bond angle, 120o: Boron and aluminium are in Group 3/13 with three outer valence electrons and can form three single covalent bonds with halogens (ignoring ionic compounds) - but no 'octet' of electrons on the central atom. e.g. gaseous boron trifluoride BF3, boron trichloride BCl3 and aluminium fluoride AlF3 (the latter is ionic in solid). Boron and aluminium are in group 3/13 with three outer shell valence electrons each of which is paired with an outer valence shell electron of the halogen. (others in the gaseous state e.g. if Q = F or Cl, then X = B, for Al for F only) but for X=Q double bonds see sulfur trioxide COCl2 (below) and the carbonate ion CO32- are also trigonal planar, no lone pairs, but involve single and double bonds Another mention of carbocations: There will always be a trigonal planar arrangement of the three C-X bonds around the positive carbon atom of a carbocation e.g. (CH3)3C+, but it will NOT be a planar molecular ion unless X is a single atom and the three Q's identical as in the case of hydrogen. Does [CCl3]+ exist?, if so, it will be trigonal planar in shape (diagram to do) |
(many other commonly used names! e.g. carbonyl chloride, carbon oxychloride, carbon dichloride oxide, phosgene-gas warfare agent). The dot and cross diagram shows this is another example of three groups of bonding electrons (two single C-Cl bonds and a C=O double bond) giving a trigonal planar structure with Cl-C-Cl and O=C-Cl bond angles of ~120o. The VSEPR theory predicts the correct shape based on three groups of bonding electrons and no lone pairs on electrons on the central carbon atom. |
Four groups of electrons around the central atom (see also sulfate and sulfite ions for four groups of bonding electrons involving single and double bonds) In all the examples in these groups there are 8 electrons in the valence shell of the central atom. |
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the Lewis dot and cross electronic diagram used to predict the shape of water, hydrogen sulfide and the amide ion - tetrahedral electron pair geometry |
![]() the 3D shape V or bent shaped molecule |
electrons: two bond pairs and two lone pairs Electron pair geometry different from the molecular geometry. (molecular compounds/ions involving hydrogen) Oxygen and sulfur in group 6/16 have six outer shell valence electrons and by pairing with two hydrogen atoms you get the stable octet shell of electrons around the central atom.
The exact H-O-H angle in H2O is reduced from 109.5o to 104.5o due to the extra repulsion of two lone pairs, the H-N-H bond angle is 107.5o in NH3 (one lone pair) and the H-C-H bond angle is 109o (no lone pairs). This trend results from the 'repulsion order' lone pair-lone pair > lone pair-bond pair > bond pair-bond pair trend - details on these other molecules later. All electrons shown on the left dot and cross diagram.
The expected bond angle would be ~109o, but, just as with the water molecule, the lone pair - lone pair repulsion is greater than the bond pair - bond pair repulsion giving a bond angle of 104.7o, almost identical with water (104.5o), and note that oxygen and nitrogen have similar atomic radii. Note for the amide ion in the o and x diagram, prior to bonding, the two black dots of the hydrogen electrons, the five black crosses of nitrogen's valence electrons and the extra purple cross of the excess electron which gives the amide group its overall single negative charge to form the amide anion. |
![]() the Lewis dot and cross electronic diagram used to predict the shape of fluorine oxide, bromine oxide, sulfur dichloride - tetrahedral electron pair geometry |
![]() V or bent shaped molecule |
electrons: two bond pairs and two lone pairs Electron pair geometry different from the molecular geometry. (molecular compounds/ions NOT involving hydrogen) Oxygen and sulfur in group 6/16 have six outer shell valence electrons and by pairing with two halogen atoms you get the stable octet shell of electrons around the central atom. The VSEPR theory argument gives a BENT or V angular shape: e.g. fluorine oxide (oxygen(II) fluoride) X = O, Q = F
The actual bond angle is 103o, and this does fit in with the argument for water where the greater lone pair - lone pair > bond pair - bond pair electron repulsion reduces the F-O-F bond angle from 109o to 103o, but read on ....!
(X = O, Q = Br) With the even more bulky bromine atom, the Br - Br repulsion completely overrides the lone pair - lone pair repulsion effect giving a Br-O-Br bond angle of 112o in Br2O (dibromine monoxide, bromine(I) oxide). Another nice pattern in a changing bond angle trend (but do pre-university exam boards realise this?.
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![]() the Lewis dot and cross electronic diagram used to predict the shape - tetrahedral electron pair geometry |
![]() pyramidal shaped molecule/ion |
electrons: three bond pairs and one lone pair Electron pair geometry different from the molecular geometry. (molecular compounds/ions involving hydrogen) Nitrogen and phosphorus in group 5/15 have five outer valence shell electrons, so when three electrons pair up with hydrogen you get three bond pairs and one lone pair and the central atom has a stable octet of electrons around it. The VSEPR theory argument gives a TRIGONAL PYRAMIDAL or TRIGONAL PYRAMID shape: e.g. ammonia NH3 with bond angle of approximately 109o. Why isn't the H-N-H angle 109o? The exact H-N-H angle is 107o due to the extra repulsion of one lone pair (see below).
and similar for the oxonium ion H3O+, H-O-H bond angle ~109o shown below where the addition of a hydrogen ion produces three bond pairs and a lone pair of electrons.
one lone pair of electrons left on protonation, the H-O-H bond angle should be <109o. Theoretically calculated value in the liquid is 106.7o, but I can't find an experimental value on the internet? This value fits in with the lone pair - bond pair repulsion > bond pair - bond pair repulsion reducing the bond angle from 109.5o, exactly as in ammonia (107.5o).
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![]() the Lewis dot and cross electronic diagram used to predict the shape - tetrahedral electron pair geometry |
![]() pyramidal shaped molecule |
electrons: three bond pairs and one lone pair Electron pair geometry different from the molecular geometry. (molecular compounds/ions NOT involving hydrogen) Nitrogen and phosphorus in group 5/15 have five outer valence shell electrons, so when three electrons pair up with halogen atoms you get three bond pairs and one lone pair and the central atom has a stable octet of electrons around it. The VSEPR theory argument gives a TRIGONAL PYRAMIDAL or TRIGONAL PYRAMID shape. e.g. nitrogen trifluoride NF3, nitrogen trichloride NCl3, or phosphorus(III) fluoride PF3, or phosphorus(III) chloride PCl3 (phosphorus trifluoride/trichloride). NF3, NCl3, PF3 and PCl3 should have bond angles Q-X-Q of approximately ~109o (Q = F, Cl etc. X = N, P etc.) The bond angles maybe <109o due to the larger lone pair - bond pair > bond pair - bond pair repulsion BUT the halogen atoms are more bulky than hydrogen. Is the bond angle trend: Br-P-Br > Cl-P-Cl > F-P-F (but I can't find actual values yet?) The Cl-N-Cl bond angle is ~107.5, similar to ammonia, presumably due to lone pair - bond pair > bond pair - bond pair repulsion, but the Cl atoms are much more bulky than H atoms, so I'm surprised it is lowered? Not sure at all on these last two points. |
![]() the Lewis dot and cross electronic diagram used to predict the shape of methane - tetrahedral electron pair geometry |
![]() tetrahedral shaped molecule/ion |
electrons: 4 bond pairs, no lone pairs Electron pair geometry same as the molecular geometry. (molecular compounds involving hydrogen) Carbon and silicon are in group 4/14 with four outer valence shell electrons, each of which can pair up with the electron of a hydrogen atom to give the stable octet arrangement around the central atom. The same applies to all the other group 4/14 elements, even if they are metallic in nature. The VSEPR theory argument gives a TETRAHEDRAL shape: e.g. methane CH4, silicon hydride SiH4 with H-X-H bond angle of the perfect 109.5o and similarly ions like the ammonium ion NH4+. Note: No lone pair, no extra repulsion, no reduction in angle, therefore perfect tetrahedral angle (for H-X-H angles: CH4 > NH3 > H2O, see below).
The ammonium ion is also a perfect tetrahedral shape, so the H-N-H bond angle is109o, identical to methane.
No lone pair of electrons left on protonation of the ammonia molecule. Prior to protonation, you have five
crosses of nitrogen's valence electrons and three dots for the bonding
hydrogen electrons. The lone pair of ammonia accepts the proton in the
formation of a dative covalent bond, BUT on proton bonding, all 4 N-H
bonds are identical, even if you represent the molecule as
H3N: |
![]() the Lewis dot and cross electronic diagram used to predict the shape of tetrachloromethane etc. - tetrahedral electron pair geometry |
![]() tetrahedral shaped molecule/ion |
electrons: 4 bond pairs, no lone pairs Electron pair geometry same as the molecular geometry. (molecular compounds/ions NOT involving hydrogen) Carbon and silicon are in group 4/14 with four outer valence shell electrons, each of which can pair up with the electron of a halogen atom to give the stable octet arrangement around the central atom. The same applies to all the other group 4/14 elements, even if they are metallic in nature.
For CCl4 in the dot and cross diagram the four crosses represent carbon's outer shell valence electrons and the four sets of seven dots represent the original chlorine outer shell valence electrons. If one or more of the atoms are not identical, then there will be small deviations from the perfect Q-X-Q bond angle of 109.5o. X can be C, Si etc. and Q can be H, F, Cl, Br etc. e.g. CF4, CBr4, CI4, SiF4, SiCl4 but also 'mixed' substituents e.g. CBr2Cl2, CHCl3, CH2F2, CBr2F2 etc. etc. !!! All of which will be nearly or perfect tetrahedral shaped molecules with all Q-X-Q bond angles of ~109.5o. All involve four pairs of bonding electrons around the central C, Si or any other group 4/14 atom,
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Five groups of electrons around the central atom |
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the Lewis dot and cross electronic diagram used to predict the shape - trigonal bipyramid electron pair geometry |
trigonal bipyramidal shaped molecule |
electrons: 5 bond pairs, no lone pairs Electron pair geometry same as the molecular geometry. Phosphorus in group 5/15 has five outer shell valence electrons, all of which can pair up with a halogen atom such as fluorine and chlorine (NOT hydrogen for 5 valency phosphorus). Note that this expands the valence shell electrons of bonded phosphorus to ten, two beyond the 'octet rule'. It will be the same for other Group 5/15 elements. The VSEPR theory argument gives a TRIGONAL BIPYRAMID shape: e.g. for phosphorus(V) fluoride (phosphorus pentafluoride) PF5, gaseous phosphorus(V) chloride, PCl5, with bond angles on 90o and 180o based on the vertical and right-angled Q-X-Q bonds and bond angles of 120o based on the central trigonal planar arrangement. In the dot and cross diagram on the left, X = P or As (but NOT N, can't use 3d level as a valence shell) and Q = F, Cl and Br (not I). Note that solid PCl5 has an ionic structure and is not a trigonal bipyramidal (bipyramid) molecule - a 4 bond pair tetrahedral [PCl4]+ ion and a 6 bond pair octahedral [PCl6]- ion, neither of which involve lone pairs. These have been separately described in the appropriate sections. |
![]() the Lewis dot and cross electronic diagram used to predict the shape - trigonal bipyramid electron pair geometry |
![]() seesaw shaped molecule/ion (if you rotate it anti-clockwise 90o to left, you can 'seesaw' up and down!) |
electrons: 4 bond pairs and one lone pair Electron pair geometry different from the molecular geometry. With the sulfur(IV) fluoride SF4 molecule, electron pairs are based on a trigonal bipyramid geometry. Two of the fluorine atoms and the lone pair are in a trigonal planar arrangement - all the F-S-F angles are 102o except the one described below. The other 2 fluorine atoms are above and below this trigonal planar arrangement with a bond angles of 173o Its described on various internet sites as a seesaw shape! The [ClF4]- and [BrF4]- ions have exactly the same dot and cross diagram, with the same shape and presumably similar bond angles. Note in all cases this expands the valence shell electrons of the central atom to ten, two beyond the 'octet rule'. |
![]() trigonal bipyramid electron pair geometry |
![]() T-shaped molecule |
electrons: 3 bond pairs and two lone pairs Electron pair geometry different from the molecular geometry. Chlorine(III) fluoride ClF3, bromine(III) fluoride BrF3 and iodine(III) fluoride IF3 are all T shape molecules (NOT a trigonal planar shape). The F-Cl-F bond angles are ~90o and ~180o. Note in all cases this expands the valence shell electrons of central halogen atom to ten, two beyond the 'octet rule'. |
![]() trigonal bipyramid electron pair geometry |
![]() linear shaped ion/molecule |
electrons: 2 bond pairs and 3 lone pairs Electron pair geometry different from the molecular geometry. The [ClF2]- ion is linear shape (contrast with the bent V shaped [ClF2]+ ion previously described). The 10 outer electrons associated with the central Cl atom comprise chlorine's original 7 out electrons, 2 from the fluorine atoms and the 10th (blue blob!) is added to give the overall single negative charge on the ion. You would expect, applying VSEPR theory, I would expect the three lone pairs to take up a trigonal planar arrangement and the 2 fluorine atoms above and below giving a F-Cl-F bond angle of 180o. All the above arguments apply to the [ICl2]- ion too - no need for extra diagrams - just swap the Cl with an I and the F's with Cl's - job done! The xenon(II) fluoride molecule XeF2 (xenon difluoride) is also a linear molecule with a F-Xe-F bond angle of 180o. Note in all cases this expands the valence shell electrons of the central atom to ten, two beyond the 'octet rule'. |
Six groups of electrons around the central atom Note in all cases here the valence shell electrons of the central atom is expanded to 12, four beyond the 'octet rule'. |
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the Lewis dot and cross electronic diagram used to predict the shape - octahedral electron pair geometry |
octahedral shaped molecule/ion The specific diagrams for the three molecules/ions mentioned are down on the right. |
electrons: 6 bond pairs, no lone pairs Electron pair geometry same as the molecular geometry. The VSEPR theory argument gives a OCTAHEDRAL SHAPE: e.g. sulfur(VI) fluoride (sulfur hexafluoride molecule) SF6 or the [PCl6]- ion and many transition metal complexes (see below), with Q-X-Q bond angles of 90o and 180o. Sulfur is in group 6/16 with six outer valence electrons and each can pair up with a halogen atom's outer electron to form six single covalent bonds. In the dot and cross diagram X = S or P and Q = F or Cl Prior to bonding, for the SF6 molecule the six dots represent the six valence shell electrons of sulfur (group 6/16) and the six sets of seven crosses represent the seven valence electrons of the fluorine atoms. Although phosphorus is in group 5/15, it can form six single covalent bonds by gaining an extra electron in the formation of a negative ion. For the [PCl6]- ion the six 'blob' dots represent the original five valence electrons of phosphorus (group 5/15) plus one extra blue dot for the extra electron that gives the ion its single overall negative charge. Similarly to above, the six sets of seven crosses represent the seven valence electrons of the chlorine atoms. You should put the dot and cross diagram in big square brackets and a superscript minus sign. The same arguments (as for P) apply to the [SiF6]2- ion, which is also octahedral - where there are two extra blue dot electrons in silicon's valence shell to give the double negative charge on the anion. Note that this expands the valence shell electrons of bonded phosphorus, silicon or sulfur to twelve, four beyond the 'octet rule'.
Note: The blue 'blob' electrons on the Si and P ox diagrams represent the extra electrons added to their valance shell to give the surplus charge of the anion.
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![]() square pyramidal shaped molecule |
electrons: 5 bond pairs, 1 lone pair Electron pair geometry different from the molecular geometry. chlorine(V) fluoride ClF5 (chlorine pentafluoride) is a square pyramid shaped molecule, with F-Cl-F bond angles of ~90o (just less than 90o?) and ~180o (actually 175o). Chlorine in group 7/17 has seven outer shell valence electrons and here five of them pair up with an outer shell electron of fluorine. The two unused electrons form the lone pair. |
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![]() square planar shaped molecule |
electrons: 4 bond pairs, 2 lone pairs Electron pair geometry different from the molecular geometry. Xenon(IV) fluoride (xenon tetrafluoride) XeF4 is a square planar shaped molecule giving a F-Xe-F bond angles of 90o and 180o. The two lone pairs will be above and below the plane of the 5 atoms of the molecule. This shape minimises the repulsion between the non-bonding pairs of electrons. Xenon in group 0/18 has eight outer shell valence electrons and here four of them pair up with an outer shell electron of fluorine. The four unused valency electrons form the two lone pairs. |
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![]() square planar shaped ion |
electrons: 4 bond pairs, 2 lone pairs Electron pair geometry different from the molecular geometry. The [ICl4]- ion is a square planar shaped molecule giving a Cl-I-Cl bond angles of 90o and 180o. The two lone pairs will be above and below the plane of the 5 atoms of the molecule. This shape minimises the repulsion between the non-bonding pairs of electrons. The [BrF4]- ion is exactly the same. Halogens in group 7/17 have seven outer shell valence electrons and here four of them pair up with an outer shell electron of another halogen. The three unused electrons, plus the one to give the overall single negative charge form the two lone pairs of non-bonding electrons. |
More complex inorganic molecules/ions and organic molecules These often are not given a particular shape name, but never-the-less, an appreciation of the 3D spatial arrangement is expected e.g. |
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See other page for more shape and bond angle analysis of organic molecules In all alkanes you are dealing with tetrahedral arrangements of bonds around the carbon atoms.
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H3N:=>BF3 would be like ethane above |
Boron trifluoride with 3 bonding pairs acts as a lone pair acceptor (Lewis acid) and ammonia with a lone pair of electrons enables it to act as a Lewis base - a an electron pair donor. It donates the lone pair to the 4th 'vacant' boron orbital to form a sort of 'adduct' compound. This gives a stable 'octet' of electrons around both the nitrogen and boron atoms. Its shape is essentially the same as ethane, a sort of double tetrahedral with H-N-H, N-B-F and F-B-F bond angles of ~109o. |
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Benzene is a completely planar molecule, with all C-C-C or C-C-H bond angles of 120o. See other page for more on shape and bond angle analysis of organic molecules |
TRANSITION METAL COMPLEX IONS
COMPLEXES The three examples below show cis/trans isomerism
more details and examples on the "Transition Metals" pages
All the bonds shown,__
or
....,
are dative covalent, with lone electron pair donation by the ligand
L, to the
central metal ion
i.e. L: APPENDIX 1 Electron pair geometry and molecular geometry The electron pair geometry is the arrangement of the (usually) pairs of bonding and non-bonding electrons around the central atom of the neutral molecule (or molecular ion). This is decided by applying the valence shell electron pair repulsion theory (VSEPR theory) to minimise the repulsion between the groups of electrons.
The molecular geometry shows the actual shape of the molecule. If there are no lone pairs of non-bonding electrons, then the electron pair geometry is the same as the molecular geometry. If there is at least one lone pair of electrons (or a single electron e.g. in a radical) then the two geometries cannot be the same.
The shapes and bond angles of BeH2 BeCl2 CO2 [Ag(NH3)2]+ BH3 BF3 BCl3 AlF3 COCl2 H2O H2S NH3 F2O PF3 PF5 PCl3 PCl5 H3O+ NCl3 CH4 CCl4 PCl4+ PCl6- SF6 H3NBF3 NH3BF3 dot and cross diagrams bond angles H-B-H VSEPR molecule shape of BH3 bond angles H-C-H VSEPR molecule shape of CH3+ bond angles F-B-F VSEPR molecule shape of BF3 bond angles Cl-B-Cl VSEPR molecule shape of BCl3 bond angles F-Al-F VSEPR molecule shape of AlF3 bond angles VSEPR molecule shape of COCl2 bond angles VSEPR molecule shape of H2S bond angles H-O-H VSEPR molecule shape of H2O bond angles H-N-H VSEPR molecule shape of NH2- bond angles VSEPR molecule shape of [NH2]- bond angles F-O-F VSEPR molecule shape of F2O bond angles Cl-O-Cl VSEPR molecule shape of Cl2O bond angles Br-O-Br VSEPR molecule shape of Br2O bond angles Cl-S-Cl VSEPR molecule shape of SCl2 bond angles F-Cl-F VSEPR molecule shape of ClF2+ bond angles VSEPR molecule shape of [ClF2]+ bond angles H-N-H VSEPR molecule shape of NH3 bond angles H-P-H VSEPR molecule shape of PH3 bond angles H-O-H VSEPR molecule shape of H3O+ bond angles VSEPR molecule shape of [H3O]+ bond angles H-C-H VSEPR molecule shape of CH3- bond angles VSEPR molecule shape of [CH3]- bond angles VSEPR molecule shape of CH3. radical bond angles Cl-N-Cl VSEPR molecule shape of NCl3 bond angles F-N-F VSEPR molecule shape of NF3 bond angles F-P-F VSEPR molecule shape of PF3 bond angles Cl-P-Cl VSEPR molecule shape of PCl3 bond angles H-C-H VSEPR molecule shape of CH4 bond angles H-Si-H VSEPR molecule shape of SiH4 bond angles H-N-H VSEPR molecule shape of NH4+ bond angles VSEPR molecule shape of [NH4]+ bond angles Cl-C-Cl VSEPR molecule shape of CCl4 bond angles F-C-F VSEPR molecule shape of CF4 bond angles Br-C-Br VSEPR molecule shape of CBr4 bond angles I-C-I VSEPR molecule shape of CI4 bond angles Cl-P-Cl VSEPR molecule shape of PCl4+ bond angles Cl-P-Cl VSEPR molecule shape of [PCl4]+ bond angles F-P-F VSEPR molecule shape of PF5 bond angles Cl-P-Cl VSEPR molecule shape of PCl5 bond angles F-S-F VSEPR molecule shape of SF6 bond angles Cl-P-Cl VSEPR molecule shape of PCl6- bond angles VSEPR molecule shape of [PCl6]- dot and cross diagrams VSEPR theory beryllium hydride BeH2 linear gaseous molecules beryllium fluoride BeF2 beryllium chloride BeCl2 beryllium bromide BeBr2 beryllium iodide BeI2 bond angle methanenitrile hydrogen cyanide HCN bent V shaped molecules trigonal planar molecules boron hydride BH3 CH3+ CCl3+ [CH3]+ [CCl3]+ carbonyl dichloride carbonyl chloride, carbon oxychloride, carbon dichloride oxide, phosgene-gas warfare agent electron pair geometry molecular geometry molecular shapes water H2O H2S bent V shaped molecules 2 bond pairs 2 lone pairs amide ion NH2- [NH2]- OF2 F2O F-O-F bond angle Br2O Br-O-Br SCl2 Cl-S-Cl sulfur dichloride sulfur(II) chloride FCl2+ ion [OF2]+ trigonal bipyramid trigonal pyramidal shape oxonium ion ammonia methyl carbanion methyl radical NF3, NCl3, PF3 PCl3 bond angles VSEPR theory tetrahedral shape CH4 SiH4 CCl4 SiCCl4 NH4+ [NH4]+ PCl4+ [PCl4]+ trigonal bipyramid trigonal bipyramidal shape PF5 PCl5 seesaw shape SF4 [ClF4]- [BrF4]- ions ClF3 BrF3 IF3 T-shaped molecules linear shaped molecules ClF2- [ClF2]- ICl2- [ICl2]- XeF2 octahedral shape SF6 PCl6- [PCl6]- SiF62- [SiF6]2- ions ClF5 square pyramid XeF4 square planar ICl4- [ICl4]- BrF4- [BrF4]-
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