What is the name of SF4?

One needs to know some basic properties of the given compound and its Lewis structure to understand its molecular geometry, polarity, and other such properties. SF4 is a chemical formula for Sulfur Tetrafluoride. It is a colorless corrosive gas that is used in the synthesis of several organofluorine compounds. SF4 is a rather hazardous compound but is used widely in chemical and pharmaceutical companies.

Table of Contents Show

SF4 Molecular Geometry
SF4 Lewis Structure
Is SF4 polar?
SF4 Hybridization
SF4 Bond angles and shape

Name of molecule	Sulfur Tetraflouride ( SF4)
No of Valence Electrons in the molecule	34
Hybridization of SF4	sp3 hybridization
Bond Angles	102 degrees and 173 degrees
Molecular Geometry of SF4	Trigonal bipyramidal

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To understand this molecule’s properties, such as its reactivity, polarity, and more, one needs to know the SF4 Lewis structure first.

SF4 Molecular Geometry

It is easy to understand the molecular geometry of a given molecule by using the molecular formula or VSEPR model. A molecular formula helps to know the exact number and type of atoms present in the given compound. Here there is one sulfur atom and four fluorine atoms in the compound, which makes it similar to the molecular formula of AX4E.

Molecules having a molecular formula of AX4E have trigonal bipyramidal molecular geometry. Here two fluorine atoms forming bonds with the sulfur atom are on the equatorial positions, and the rest two are on the axial positions. As there is one lone pair on the central atom, it repels the bonding pair of electrons, which tweaks the shape a little bit and makes it appear like a see-saw. The electrons follow this pattern of arrangement following the VSEPR rule to minimize the repulsion forces between the lone pairs of electrons to maximize the molecule’s stability.

Hence, SF4 has a trigonal bipyramidal molecular geometry.

SF4 Lewis Structure

Lewis structure is a pictorial representation of the bonds and valence electrons in the molecule. The bonds formed between two atoms are depicted using lines, whereas the valence electrons not forming any bonds are shown by dots. The valence electrons that participate in forming bonds are called bonding pairs of electrons, whereas the electrons that do not participate or form any bonds are called nonbonding pairs of electrons or lone pairs.

And to draw the Lewis structure of SF4, we first need to know the total number of valence electrons in this molecule.

As one can probably see, there is one sulfur atom in this compound and four fluorine atoms. To know the total valence electrons of this compound, we need to know the valence electrons of both the atoms individually.

Valence electrons of Sulfur: 6
Valence electrons of Fluorine: 4* (7)

( as there are four fluorine atoms, we have to consider valence electrons of all atoms)

Total number of valence electrons in SF4 = number of valence electrons in sulfur + number of valence electrons in fluorine

= 6 + 28

= 34 valence electrons

Now that we know the total number of valence electrons, it would become easy for us to understand the bond formation between the atoms and the complete arrangement of the molecule too.

Sulfur will be the central atom in this molecule as it is the least electronegative, with four fluorine atoms forming bonds on the sides of this central atom. Every fluorine atom will form a bond with the central atom, which means there will be four bonds in the molecule structure using up four valence electrons of fluorine atoms and 4 electrons of the sulfur atom. So now, eight valence electrons are used, reducing the number of valence electrons from 34 to 24. All the fluorine atoms have six valence electrons, and the central atom has two valence electrons.

Draw lines between S and F to show bonds and for lone pairs of electrons, use dots. Each fluorine atom will have three pairs of 6 valence electrons ( shown as dots) on the atom, along with one bond with sulfur. In contrast, the central atom will have two valence electrons and four bonds.

Hence, the central atom, sulfur, will have one lone pair of electrons and four bonding pairs of electrons in the Lewis structure of SF4. At the same time, each fluorine atom will have three lone pairs.

Is SF4 polar?

Once we know the Lewis structure and molecular geometry of the given compound, it becomes easier to depict the molecule’s polarity. Here, one lone pair on the central sulfur atom and four bonding pairs of electrons leads to the asymmetric distribution of electrons on the central atom.

Also, as the shape of the molecule is like a see-saw, two fluorine atoms can cancel out each other’s dipole moment, but the rest two can’t due to the electrons’ arrangement. And as fluorine atoms are more electronegative than the sulfur atom, it results in uneven distribution of the charge. Hence the dipole moment is not canceled, which makes the molecule polar. So yes, SF4 is polar.

SF4 Hybridization

To know the hybridization of the SF4 molecule, let us first look at the regions of electron density for the central atom.

Sulfur has four bonding pairs of electrons and one lone pair, making its total number of regions for electron density 5. Hence the sulfur atom uses five hybridized orbitals, one 3s orbital, three 3p orbitals, and one 3d orbital. This arrangement of electrons around the atom and hybridized orbitals leads to the sp3d hybridization. One can also use the steric number to know the hybridization; here, the steric number is 5 for the sulfur atom.

Thus SF4 has sp3d hybridization.

SF4 Bond angles and shape

The central sulfur atom forms four bonds with the neighboring fluorine atoms and has one lone pair of electrons. Fluorine atoms on the equatorial positions have the bond angles of 102 degrees, and the axial ones have 173 degrees, which are a little different than the trigonal bipyramidal molecular geometry leading to a see-saw shape.

The lone pair on the central atom leads to the change in the bond angles from 120 degrees to 102 degrees for equatorial fluorine atoms and 173 degrees instead of 180 degrees for axial fluorine atoms.

Concluding Remarks

To conclude all the properties we can say that,

Sulfur Tetrafluoride has 34 valence electrons, out of which it forms four covalent bonds and one lone pair of electrons on the central atom in its Lewis structure.
There are three lone pairs on each fluorine atom.
It has a molecular geometry of the formula AX4E; it forms a see-saw shape and has a trigonal bipyramidal molecular geometry.
SF4 has sp3d hybridization and is polar in nature.

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Sulfur tetrafluoride is the chemical compound with the formula SF4. It is a colorless corrosive gas that releases dangerous HF upon exposure to water or moisture. Despite these unwelcome characteristics, this compound is a useful reagent for the preparation of organofluorine compounds,[3] some of which are important in the pharmaceutical and specialty chemical industries.

Sulfur tetrafluoride

Names IUPAC name

Sulfur(IV) fluoride

Other names

Sulfur tetrafluoride

Identifiers

CAS Number

7783-60-0
Y

3D model (JSmol)

Interactive image

ChEBI

CHEBI:30495
Y

ChemSpider

22961
Y

ECHA InfoCard 100.029.103

PubChem CID

24555

RTECS number

WT4800000

UNII

F4P8J39GOF
Y

UN number 2418

CompTox Dashboard (EPA)

DTXSID40893074

InChI

InChI=1S/F4S/c1-5(2,3)4
Y
Key: QHMQWEPBXSHHLH-UHFFFAOYSA-N
Y
InChI=1/F4S/c1-5(2,3)4
Key: QHMQWEPBXSHHLH-UHFFFAOYAT

SMILES

FS(F)(F)F

Properties

Chemical formula

SF4 Molar mass 108.07 g/mol Appearance colorless gas Density 1.95 g/cm3, −78 °C Melting point −121.0 °C Boiling point −38 °C

Solubility in water

reacts Vapor pressure 10.5 atm (22°C)[1]Structure

Molecular shape

Seesaw (C2v)

Dipole moment

0.632 D[2]Hazards Occupational safety and health (OHS/OSH):

Main hazards

highly toxic
corrosive NFPA 704 (fire diamond)

NIOSH (US health exposure limits):

PEL (Permissible)

none[1]

REL (Recommended)

C 0.1 ppm (0.4 mg/m3)[1]

IDLH (Immediate danger)

N.D.[1]Safety data sheet (SDS) ICSC 1456 Related compounds

Other anions

Sulfur dichloride
Disulfur dibromide
Sulfur trifluoride

Other cations

Oxygen difluoride
Selenium tetrafluoride
Tellurium tetrafluoride

Related sulfur fluorides

Disulfur difluoride
Sulfur difluoride
Disulfur decafluoride
Sulfur hexafluoride

Related compounds

Thionyl fluoride

Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Y verify (what is

N ?)

Infobox references

Sulfur in SF4 is in the formal +4 oxidation state. Of sulfur's total of six valence electrons, two form a lone pair. The structure of SF4 can therefore be anticipated using the principles of VSEPR theory: it is a see-saw shape, with S at the center. One of the three equatorial positions is occupied by a nonbonding lone pair of electrons. Consequently, the molecule has two distinct types of F ligands, two axial and two equatorial. The relevant bond distances are S–Fax = 164.3 pm and S–Feq = 154.2 pm. It is typical for the axial ligands in hypervalent molecules to be bonded less strongly. In contrast to SF4, the related molecule SF6 has sulfur in the 6+ state, no valence electrons remain nonbonding on sulfur, hence the molecule adopts a highly symmetrical octahedral structure. Further contrasting with SF4, SF6 is extraordinarily inert chemically.

The 19F NMR spectrum of SF4 reveals only one signal, which indicates that the axial and equatorial F atom positions rapidly interconvert via pseudorotation.[4]

Intramolecular dynamic equilibration of SF4.

SF4 is produced by the reaction of SCl2 and NaF in acetonitrile:[5]

3 SCl2 + 4 NaF → SF4 + S2Cl2 + 4 NaCl

SF4 is also produced in the absence of solvent at elevated temperatures.[6][7]

Alternatively, SF4 at high yield is produced using sulfur (S), NaF and chlorine (Cl2) in the absence of reaction medium, also at less-desirable elevated reaction temperatures (e.g. 225–450 °C).[6][7]

A low temperature (e.g. 20–86 °C) method of producing SF4 at high yield, without the requirement for reaction medium, has been demonstrated utilizing bromine (Br2) instead of chlorine (Cl2), S and KF:[8]

S + (2 + x) Br2 + 4 KF → SF4↑ + x Br2 + 4 KBr

In organic synthesis, SF4 is used to convert COH and C=O groups into CF and CF2 groups, respectively.[9] Certain alcohols readily give the corresponding fluorocarbon. Ketones and aldehydes give geminal difluorides. The presence of protons alpha to the carbonyl leads to side reactions and diminished (30–40%) yield. Also diols can give cyclic sulfite esters, (RO)2SO. Carboxylic acids convert to trifluoromethyl derivatives. For example, treatment of heptanoic acid with SF4 at 100–130 °C produces 1,1,1-trifluoroheptane. Hexafluoro-2-butyne can be similarly produced from acetylenedicarboxylic acid. The coproducts from these fluorinations, including unreacted SF4 together with SOF2 and SO2, are toxic but can be neutralized by their treatment with aqueous KOH.

The use of SF4 is being superseded in recent years by the more conveniently handled diethylaminosulfur trifluoride, Et2NSF3, "DAST", where Et = CH3CH2.[10] This reagent is prepared from SF4:[11]

SF4 + Me3SiNEt2 → Et2NSF3 + Me3SiF

Sulfur chloride pentafluoride (SF
5Cl), a useful source of the SF5 group, is prepared from SF4.[12]

Hydrolysis of SF4 gives sulfur dioxide:[13]

SF4 + 2 H2O → SO2 + 4 HF

This reaction proceeds via the intermediacy of thionyl fluoride, which usually does not interfere with the use of SF4 as a reagent.[5]

SF
4 reacts inside the lungs with moisture, generating sulfur dioxide and hydrogen fluoride:[14]

SF4 + 2 H2O → SO2 + 4 HF

^ a b c d NIOSH Pocket Guide to Chemical Hazards. "#0580". National Institute for Occupational Safety and Health (NIOSH).
^ Tolles, W. M.; W. M. Gwinn, W. D. (1962). "Structure and Dipole Moment for SF4". J. Chem. Phys. 36 (5): 1119–1121. Bibcode:1962JChPh..36.1119T. doi:10.1063/1.1732702.
^ Wang, C.-L. J. (2004). "Sulfur Tetrafluoride". In Paquette, L. (ed.). Encyclopedia of Reagents for Organic Synthesis. New York: J. Wiley & Sons. doi:10.1002/047084289X. hdl:10261/236866. ISBN 9780471936237.
^ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5.
^ a b Fawcett, F. S.; Tullock, C. W. (1963). Sulfur (IV) Fluoride: (Sulfur Tetrafluoride). Inorganic Syntheses. Vol. 7. pp. 119–124. doi:10.1002/9780470132388.ch33.
^ a b Tullock, C. W.; Fawcett, F. S.; Smith, W. C.; Coffman, D. D. (1960). "The Chemistry of Sulfur Tetrafluoride. I. The Synthesis of Sulfur Tetrafluoride". J. Am. Chem. Soc. 82 (3): 539–542. doi:10.1021/ja01488a011.
^ a b US 2992073, Tullock, C.W., "Synthesis of Sulfur Tetrafluoride", issued 1961
^ Winter, R.W.; Cook P.W. (2010). "A simplified and efficient bromine-facilitated SF4-preparation method". J. Fluorine Chem. 131: 780-783. doi:10.1016/j.jfluchem.2010.03.016
^ Hasek, W. R. "1,1,1-Trifluoroheptane". Organic Syntheses.; Collective Volume, vol. 5, p. 1082
^ Fauq, A. H. (2004). "N,N-Diethylaminosulfur Trifluoride". In Paquette, L. (ed.). Encyclopedia of Reagents for Organic Synthesis. New York: J. Wiley & Sons. doi:10.1002/047084289X. hdl:10261/236866. ISBN 9780471936237..
^ W. J. Middleton; E. M. Bingham. "Diethylaminosulfur Trifluoride". Organic Syntheses.; Collective Volume, vol. 6, p. 440
^ Nyman, F.; Roberts, H. L.; Seaton, T. (1966). "Sulfur Chloride Pentafluoride". Inorganic Syntheses. Inorganic Syntheses. Vol. 8. McGraw-Hill. p. 160. doi:10.1002/9780470132395.ch42. ISBN 9780470132395.
^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
^ Johnston, H. (2003). A Bridge not Attacked: Chemical Warfare Civilian Research During World War II. World Scientific. pp. 33–36. ISBN 981-238-153-8.