(c) doc bDoc 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

The shapes and bond angles of a variety of molecules are described, explained and discussed using valence shell electron pair repulsion theory (VSEPR theory) and patterns of shapes deduced for 2, 3, 4, 5 and 6 groups of bonding electrons or non-bonding electrons in the valence shell of the central atom of the molecule or ion. 

So, this page is all about how to work out molecule shapes and work out bond angles is described and explained! i.e. using the electron pair repulsion theory to deduce the shapes of molecules and work out bond angles from the various repulsions i.e lone pair - lone pair, lone pair - bond pair and bond pair - bond pair repulsions. Shapes can be worked out from dot & cross diagrams and bond angles deduced from the established shape, all such deductions are essentially based on electron pair(s) repulsion theory. A summary of the shapes of molecules is given in APPENDIX 1.

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 * No-shape diatomic molecules!

Molecule shapes based on total groups of 2  3  4  5  6  electron pairs (bonding or non-bonding)

Some more complex inorganic/organic molecules/ions

transition metal complexes

Some other molecules/ions of carbon, nitrogen, sulfur and chlorine

shapes and bond angles of organic molecules

GCSE/IGCSE/AS INDEX of Science-CHEMISTRY bonding notes

 


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.

  1. Consider bonding pairs and lone (non-bonding) pairs of electrons as charge clouds that repel each other (remember that like electrical charges repel).
    • In deducing the shape of free radicals, a single electron can be treated as a lone pair to a reasonable approximation.
  2. These pairs of bonding electrons or non-bonding electrons in the outer shell of atoms arrange themselves as far apart as possible to minimise this repulsion.
    • We are talking about the electron clouds or more precisely, the orbitals that these bonding and non-bonding pairs of electrons occupy.
    • This effectively means to produce a range of as wide as possible angles between adjacent orbitals.
    • These maybe lone pair - lone pair, bond pair - lone pair or bond pair - bond pairs of electrons.
    • Note that the electron cloud orbitals could involve double or triple bonds - but that makes no difference in deducing the shape of a molecule using VSEPR theory.
  3. Know the general rule: lone pair–lone pair repulsion is greater than lone pair–bond pair repulsion, which is greater than bond pair–bond pair repulsion.
    • Generally speaking this doesn't affect the shape, BUT, it can make small differences to the expected 'perfect' bond angle. An excellent example is the comparison of the H-X-H angle for water, ammonia and methane, which I've discussed in detail on this page.
    • You also have to be careful in predicting bond angles using the rule expressed in 2. which works well if X is nitrogen or oxygen.
    • BUT, larger X atoms or more bulky X groups might override the valence shell electron pair rule of 2.
      • This can be overlooked by exam boards in their marking schemes and 'stuff' you find on the internet!
    • Frequently in the text I've used the abbreviated term 'lone pair' meaning a pair of non-bonding electrons - in exams its a good idea to initially define any abbreviations you may use several times and then you shouldn't lose any marks - take care!
  4. Last introduction point: Be very careful when describing your deduction of a molecular shape to distinguish between the electron pair geometry and the actual molecular geometry of the molecule i.e. the actual shape! If no lone pairs are involved, they are the same, if at least one lone pair is involved, they cannot be the same!
    • 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).
    • The molecular geometry shows the actual shape of the molecule.
    • Check out the extensive summary table in APPENDIX 1, it covers all shapes needed for UK A level chemistry.
    • I haven't introduced this APPENDIX 1 summary here at the start - its too much in one go, but when you have studied a variety of examples of deducing molecule shapes and bond angles take a good look at it and you should then fully appreciate point 4.
 
Introduction continued - 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.

As already mentioned, this has an important 'sub-rule' which affects the precise bond angle.

Any lone pairs of non-bonding electrons on the central atom X, are closer to X than bond pairs because there is no Q atom attracting/sharing the lone pair electron charge.

This will increase the repulsion between a lone pair of non-bonding electrons on X and any other bonding/non-bonding pairs of electrons in the outer shell of the central atom X.

The result is two-fold:

In terms of electron pair repulsion: lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.

As the lone pair - 'other pair' repulsion increases, the angle between these pairs increases, so the Q-X-Q angle will be slightly reduced compared to what might be expected from the 'simple' geometry of the shape (this is best illustrated by the sequence H2O, NH3 and CH4, see below)

 

 

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.

(c) doc b 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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(c) doc b Cl-Cl e.g. chlorine Cl2, iodine(I) chloride ICl (iodine monochloride)

(c) doc b Q and X are both halogen atoms from group 7

  O=O (c) doc b oxygen molecule

  

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)

(c) doc bLewis dot and cross electronic diagram used to predict the shape of beryllium chloride

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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

(c) doc b carbon dioxide CO2

full dot and cross electron structure

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-CN 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.

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the Lewis dot and cross electronic diagram used to predict the shape - trigonal planar electron pair geometry

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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.

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the Lewis dot and cross electronic diagram used to predict the shape of boron hydride - trigonal planar electron pair geometry

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trigonal planar shaped molecule

electrons: 3 bond pairs, no lone pairs

Electron pair geometry same as the molecular geometry.

(molecular compounds/ions involving hydrogen)

(c) doc bThe VSEPR theory argument gives the shape TRIGONAL PLANAR: Q-X-Q bond angle exactly 120o:

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.

(X = C, Q = H) Carbon in group 4/14 has four valence electrons, one of which is lost in the formation of carbocation (carbonium ion). When the remaining three electrons are paired with those from hydrogens you get the methyl carbocation CH3+. This will also have a perfect trigonal planar shape with a perfect H-C-H bond angle of 120o. Prior to bonding, the black dots of the hydrogen electrons and the black crosses of the carbon electrons and not a black cross has gone from the carbon atom, one of its four outer valence electrons has been lost to give he overall single positive charge.

DO NOT assume examination questions will always be based on neutral molecules!

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the Lewis dot and cross electronic diagram used to predict the shape of some group III trihalides - trigonal planar electron pair geometry

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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)

COCl2 (g) carbonyl dichloride

(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

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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.

(c) doc b The VSEPR theory argument gives an ANGULAR or BENT or V shape: e.g. hydrogen sulfide, H2S, or water H2O, i.e. H2X with H-X-H bond angle of approximately 109o (actually 104.5o for water) and similarly ions like the amide ion NH2- (Q = H, X = O, S etc. in group 6)

(c) doc b  (c) doc b The two pairs of double dots represent the lone pairs of non-bonding electrons not involved in the covalent bonding in water. Similar diagram for hydrogen sulfide H2S.

Why isn't the H-O-H angle 109o?  

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.

You would expect hydrogen sulfide H2S, to have the same bent V shape and a H-S-H bond angle of ~109o, but due to the greater repulsion of the non-bonding pairs of electrons (lone pairs) the actual H-S-H bond angle is reduced to 92o. (this great reduction from 109o, isn't just about the repulsion rule, but also involves the subtle bonding behaviour of the s and p orbitals - of no concern to us here at pre-university level!)

An example of an ion with a bent V shape is the negative ion, the amide ion HN2-.

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.

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the Lewis dot and cross electronic diagram used to predict the shape of fluorine oxide, bromine oxide, sulfur dichloride - tetrahedral electron pair geometry

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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

(c) doc bOF2 (F2O) with F-O-F bond angle of approximately 109o

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 ....!

(c) doc b(X = O, Q = Cl) Cl2O [Chlorine (I) oxide, dichlorine monoxide] is also V or bent shaped with a Cl-O-Cl bond angle of ~111o, close to the perfect 'tetrahedral' bond angle for 4 pairs of valence shell electrons. Although lone pair - lone pair > lone pair - bond pair > bond pair - bond pair repulsion, do NOT assume, as with water, the bond angle is <109o. The reason being that chlorine is a much more bulky atom than hydrogen (radii Cl >> H) and so the Cl - Cl electron cloud repulsion compensates for the extra lone pair - lone pair repulsion effect and the bond angle is actually ~110o, very close to the perfect 109.5o, but this is by coincidence! Lots of X2O molecules on the internet are quoted as having a bond angle of 104.5o, but unless X is hydrogen, they are usually wrong!

(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?.

(c) doc b(X = S, Q = Cl) Sulfur(II) chloride (sulphur dichloride) SCl2, will have the same dot and cross diagram (two bond pairs and two lone pairs of electrons) as OF2 (F2O) and Cl2O and a predicted bent or V shaped molecule. The predicted Cl-S-Cl bond angle would be ~109o. but the extra lone pair - lone pair repulsion lowers this to 103o. SF2 is similar with a bond angle of 98o but this greatly lowered angle is not simply to do with the effect of lone pair - lone pair repulsion. There are complications due to differences in electronegativity and the way the valence orbitals are distributed and discussion of these points is well beyond pre-university chemistry!

[(c) doc b]+ is the outer electronic structure of the [F-Cl-F]+ ion.

Chlorine is in group 7/17 with seven outer valence shell electrons. Remove an electron and pair up two of the remaining six electrons with fluorine and you get a dot and cross diagram with two bond pairs and two pairs of non-bonding electrons. e.g. the difluorochlorinium(III) cation [ClF2]+ ion is also bent shaped. You expect the bond angle to be ~109.

I've seen values of the F-Cl-F bond angle of 104.5 quoted on the internet as is if this ion is analogous to water? not sure, because F atoms are more bulky than H atoms (bigger group of electron clouds) giving increased F - F repulsion. I've also seen in a (Salters A level?) marking scheme 109o not allowed but 107-108o allowed. BUT, it might well be ~109o because the F - F repulsion might compensate for the extra lone pair - lone pair repulsion OR it might be ~104.5o like F2O but I can't actually find an authentic value on the internet for the F-Cl-F bond angle?

Can anybody help on this one with a clear argument for a particular value?

I've found one quote of ~109o from Google books, but this maybe just an assumption

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the Lewis dot and cross electronic diagram used to predict the shape - tetrahedral electron pair geometry

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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).

(c) doc b (c) doc b The double dots represent a lone pair of non-bonding electrons not involved in the covalent bonding in ammonia.

Note: the exact H-O-H angle in H2O is 104.5o due to the extra repulsion of two lone pairs, the H-N-H is 107.5o in NH3 (one lone pair) and H-C-H is 109o (no lone pairs) because of the 'repulsion order' lone pair - bond pair > bond pair - bond pair repulsion,

 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).

(c) doc b The shape of phosphine PH3, would be expected to be pyramidal (trigonal pyramid) like ammonia with a bond angle of ~109., BUT, in fact the experimental H-P-H bond angle is 93.5o, but this is only partly explained by the lone pair - bond pair repulsion > bond pair - bond pair repulsion, and, as with hydrogen sulfide, it probably involves the subtle bonding behaviour of the s and p orbitals (of no concern to us here at pre-university level!)

The methyl carbanion ion CH3- (negative methyl ion) will also have a trigonal pyramid shape based on three bond pairs and a lone pair. Prior to bonding the 4 black dots represent carbon's valence electrons and the crosses the electrons from the hydrogen atoms. Then, the purple dot represents the extra electron added to give the overall single negatively charged ion.

The methyl radical CH3. is also a pyramidal shape because the lone electron effectively acts as a 4th electron cloud (orbital) along with three bond pairs associated with carbon's valence shell electrons combining with the hydrogen electrons.  Prior to bonding the 4 black dots represent carbon's valence electrons and the crosses the electrons from the hydrogen atoms. With only 3 bond pairs, this leaves a lone electron in its own orbital of the electrically neutral free radical.

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the Lewis dot and cross electronic diagram used to predict the shape - tetrahedral electron pair geometry

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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.

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the Lewis dot and cross electronic diagram used to predict the shape of methane - tetrahedral electron pair geometry

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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).

(c) doc b alkanes structure and naming (c) doc b (Q = H, X = C, Si, Ge etc. in group 4)

Note: the exact H-O-H angle in H2O is 104.5o due to the extra repulsion of two lone pairs, the H-N-H is 107.5o in NH3 (one lone pair) and H-C-H is the perfect tetrahedral angle of 109.5o (no lone pairs) because of the 'repulsion order' lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.

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:H !!

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the Lewis dot and cross electronic diagram used to predict the shape of tetrachloromethane etc. - tetrahedral electron pair geometry

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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.

The VSEPR theory argument gives a TETRAHEDRAL shape: e.g. tetrachloromethane CCl4 or [PCl4]+ with exact Cl-C-Cl and Cl-P-Cl bond angles of 109.5o.

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,

You can get a tetrahedral molecule-ion from a group 5/15 element, if one of the five outer electrons is effectively lost to make a positive ion.  In the case of the [PCl4]+ ion in the dot and cross diagram the four crosses represent the remaining phosphorus outer shell valence electrons. P is in group 5/15 with five outer electrons, BUT a 'dot' has been removed from the dot and cross Lewis diagram to give the single positive charge on the ion. The four sets of seven dots represent the original chlorine outer shell valence electrons as in CCl4 above. The phosphorous supplies the four electrons needed to produce the four P-Cl covalent bonds.

 

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

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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'.

(c) doc b

the Lewis dot and cross electronic diagram used to predict the shape - octahedral electron pair geometry

(c) doc b

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'.

SiF6, [PCl6]-, [SiF6]2-

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.

SF6, [PCl6]-, [SiF6]2-

- octahedral electron pair geometry

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.

- octahedral electron pair geometry

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.

- octahedral electron pair geometry

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.

(c) doc b

Ethane consists of two joined 'tetrahedral halves', with all C-C-H and H-C-H bond angles of 109o.

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.

 

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.

(c) doc b or (c) doc b

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(c) doc b

The three examples below show cis/trans isomerism

(c) doc bcis/trans octahedral

(c) doc bcis/trans octahedral

(c) doc b cis/trans square planar

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: (c) doc b Mn+ etc.


 

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 groups of electron pairs is dealt with here, but the 'group' could be lone electron of 4 or 6 electrons for a double or triple covalent bond. Therefore, when applying VSEPR to molecules or ions with multiple bonds, you treat double or triple bonds as if they were single bonds to deduce the molecular geometry.

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.

Bond pairs around the central atom

Lone pairs of electrons around the 'central' atom

Total groups of electrons - here just pairs considered

Electron-pair geometry - considering all the outer shell pairs of electrons

Molecular Geometry - actual shape of molecule

Bond
angles o

Examples of molecular shape

The Lewis dot & cross electronic diagram and molecule shape

2

0

2

linear

linear

180

 (c) doc b Cl-Be-Cl, BeCl2

3

0

3

trigonal planar

trigonal planar

120

 (c) doc b (c) doc b BF3, BCl3

2

1

3

trigonal planar

V or bent

~ 120

:SO2, O=S=O, 3 groups of electrons BUT actually two double bonds and one lone pair of electrons

4

0

4

tetrahedral

tetrahedral

109.5

  (c) doc b CCl4

3

1

4

tetrahedral

trigonal pyramid

~109.5

can be more or less

 (c) doc b (c) doc b NH3

2

2

4

tetrahedral

V or bent

less than 109.5

can be more or less

 (c) doc b (c) doc b H2O

5

0

5

trigonal bipyramidal

trigonal bipyramid

90, 120 and 180

 (c) doc b (c) doc b PF5

4

1

5

trigonal bipyramidal

seesaw

>90 and <180?

SF4, [BrF4]+, NOT needed for A level?

3

2

5

trigonal bipyramidal

T-shaped

~90 and ~180

  ClF3

2

3

5

trigonal bipyramidal

linear

180

  [ClF2]-

6

0

6

octahedral

octahedral

90 and 180

 (c) doc b (c) doc b SF6

5

1

6

octahedral

square pyramid

~90 and ~180

  ClF5

4

2

6

octahedral

square planar

90 and 180

  XeF4

 


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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