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.

• The Number Of Valence Electrons In Fluorine
• Amount Of Valence Electrons In Fluorine Atom
• Amount Of Valence Electrons In Fluorine
• Number Of Valence Electrons In Chlorine
• Number Of Valence Electrons In Fluorine
• Number Of Valence Electrons In Fluoride
• Determine The Number Of Valence Electrons In Fluorine Atom

The electrons like to be in separate shells/orbitals. Shell number one can only hold 2 electrons, shell two can hold 8, and for the first eighteen elements shell three can hold a maximum of eight electrons. As you learn about elements with more than eighteen electrons you will find that shell three can hold more than eight. Two fluorine atoms, for example, can form a stable F 2 molecule in which each atom has an octet of valence electrons by sharing a pair of electrons. A pair of oxygen atoms can form an O 2 molecule in which each atom has a total of eight valence electrons by sharing two pairs of electrons. Total Number of Electrons in Neutral Atom Valence Electrons Gain or Lose Electrons Ion Formed Hydrogen 1 1 Gain or Lose 1 H+ or H-Helium 2 2 None None Lithium 3 1 Lose 1 Li+ Beryllium 4 2 Lose 2 Be2+ Boron 5 3 Lose 3 B3+ Carbon 6 4 Gain or Lose 4 C4+or C4-Nitrogen 7 5 Gain 3 N3-Oxygen 8 6 Gain 2 O2-Fluorine 9 7 Gain 1 F. The fluorine atom has seven valence electrons. The atomic number of fluorine is 9, and its electronic configuration is He 2s2 2p5. The atom of fluorine prefers to gain an electron to complete its last orbit. Find out the total number of valence electrons in PF3, which is 26. Find out the number of valence electrons further needed of a single PF3 molecule to stabilize itself. It is six in total, where three valence electrons are needed by the phosphorus atom and one, each by three fluorine atoms.

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.

The Number Of Valence Electrons In Fluorine

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.

Learning Objectives

• Describe and explain the observed trends in atomic size, ionization energy, and electron affinity of the elements

The elements in groups (vertical columns) of the periodic table exhibit similar chemical behavior. This similarity occurs because the members of a group have the same number and distribution of electrons in their valence shells. However, there are also other patterns in chemical properties on the periodic table. For example, as we move down a group, the metallic character of the atoms increases. Oxygen, at the top of Group 16 (6A), is a colorless gas; in the middle of the group, selenium is a semiconducting solid; and, toward the bottom, polonium is a silver-grey solid that conducts electricity.

As we go across a period from left to right, we add a proton to the nucleus and an electron to the valence shell with each successive element. As we go down the elements in a group, the number of electrons in the valence shell remains constant, but the principal quantum number increases by one each time. An understanding of the electronic structure of the elements allows us to examine some of the properties that govern their chemical behavior. These properties vary periodically as the electronic structure of the elements changes. They are (1) size (radius) of atoms and ions, (2) ionization energies, and (3) electron affinities.

Variation in Covalent Radius

The quantum mechanical picture makes it difficult to establish a definite size of an atom. However, there are several practical ways to define the radius of atoms and, thus, to determine their relative sizes that give roughly similar values. We will use the covalent radius (Figure (PageIndex{1})), which is defined as one-half the distance between the nuclei of two identical atoms when they are joined by a covalent bond (this measurement is possible because atoms within molecules still retain much of their atomic identity).

We know that as we scan down a group, the principal quantum number, n, increases by one for each element. Thus, the electrons are being added to a region of space that is increasingly distant from the nucleus. Consequently, the size of the atom (and its covalent radius) must increase as we increase the distance of the outermost electrons from the nucleus. This trend is illustrated for the covalent radii of the halogens in Table (PageIndex{1}) and Figure (PageIndex{1}). The trends for the entire periodic table can be seen in Figure (PageIndex{2}).

As shown in Figure (PageIndex{2}), as we move across a period from left to right, we generally find that each element has a smaller covalent radius than the element preceding it. This might seem counterintuitive because it implies that atoms with more electrons have a smaller atomic radius. This can be explained with the concept of effective nuclear charge, (Z_{eff}). This is the pull exerted on a specific electron by the nucleus, taking into account any electron–electron repulsions. For hydrogen, there is only one electron and so the nuclear charge (Z) and the effective nuclear charge (Z_eff) are equal. For all other atoms, the inner electrons partially shield the outer electrons from the pull of the nucleus, and thus:

[Z_ce{eff}=Z−shielding]

Shielding is determined by the probability of another electron being between the electron of interest and the nucleus, as well as by the electron–electron repulsions the electron of interest encounters. Core electrons are adept at shielding, while electrons in the same valence shell do not block the nuclear attraction experienced by each other as efficiently. Thus, each time we move from one element to the next across a period, Z increases by one, but the shielding increases only slightly. Thus, Z_eff increases as we move from left to right across a period. The stronger pull (higher effective nuclear charge) experienced by electrons on the right side of the periodic table draws them closer to the nucleus, making the covalent radii smaller.

Thus, as we would expect, the outermost or valence electrons are easiest to remove because they have the highest energies, are shielded more, and are farthest from the nucleus. As a general rule, when the representative elements form cations, they do so by the loss of the ns or np electrons that were added last in the Aufbau process. The transition elements, on the other hand, lose the ns electrons before they begin to lose the (n – 1)d electrons, even though the ns electrons are added first, according to the Aufbau principle.

Example (PageIndex{1}): Sorting Atomic Radii

Predict the order of increasing covalent radius for Ge, Fl, Br, Kr.

Solution

Radius increases as we move down a group, so Ge < Fl (Note: Fl is the symbol for flerovium, element 114, NOT fluorine). Radius decreases as we move across a period, so Kr < Br < Ge. Putting the trends together, we obtain Kr < Br < Ge < Fl.

Amount Of Valence Electrons In Fluorine Atom

Exercise (PageIndex{1})

Give an example of an atom whose size is smaller than fluorine.

Ne or He

Variation in Ionic Radii

Ionic radius is the measure used to describe the size of an ion. A cation always has fewer electrons and the same number of protons as the parent atom; it is smaller than the atom from which it is derived (Figure (PageIndex{3})). For example, the covalent radius of an aluminum atom (1s²2s²2p⁶3s²3p¹) is 118 pm, whereas the ionic radius of an Al³⁺ (1s²2s²2p⁶) is 68 pm. As electrons are removed from the outer valence shell, the remaining core electrons occupying smaller shells experience a greater effective nuclear charge Z_eff (as discussed) and are drawn even closer to the nucleus.

Cations with larger charges are smaller than cations with smaller charges (e.g., V²⁺ has an ionic radius of 79 pm, while that of V³⁺ is 64 pm). Proceeding down the groups of the periodic table, we find that cations of successive elements with the same charge generally have larger radii, corresponding to an increase in the principal quantum number, n.

An anion (negative ion) is formed by the addition of one or more electrons to the valence shell of an atom. This results in a greater repulsion among the electrons and a decrease in (Z_{eff}) per electron. Both effects (the increased number of electrons and the decreased Z_eff) cause the radius of an anion to be larger than that of the parent atom ( Figure (PageIndex{3})). For example, a sulfur atom ([Ne]3s²3p⁴) has a covalent radius of 104 pm, whereas the ionic radius of the sulfide anion ([Ne]3s²3p⁶) is 170 pm. For consecutive elements proceeding down any group, anions have larger principal quantum numbers and, thus, larger radii.

Atoms and ions that have the same electron configuration are said to be isoelectronic. Examples of isoelectronic species are N^3–, O^2–, F^–, Ne, Na⁺, Mg²⁺, and Al³⁺ (1s²2s²2p⁶). Another isoelectronic series is P^3–, S^2–, Cl^–, Ar, K⁺, Ca²⁺, and Sc³⁺ ([Ne]3s²3p⁶). For atoms or ions that are isoelectronic, the number of protons determines the size. The greater the nuclear charge, the smaller the radius in a series of isoelectronic ions and atoms.

Variation in Ionization Energies

The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy (IE₁). The first ionization energy for an element, X, is the energy required to form a cation with +1 charge:

The energy required to remove the second most loosely bound electron is called the second ionization energy (IE₂).

The energy required to remove the third electron is the third ionization energy, and so on. Energy is always required to remove electrons from atoms or ions, so ionization processes are endothermic and IE values are always positive. For larger atoms, the most loosely bound electron is located farther from the nucleus and so is easier to remove. Thus, as size (atomic radius) increases, the ionization energy should decrease. Relating this logic to what we have just learned about radii, we would expect first ionization energies to decrease down a group and to increase across a period.

Figure (PageIndex{4}) graphs the relationship between the first ionization energy and the atomic number of several elements. Within a period, the values of first ionization energy for the elements (IE₁) generally increases with increasing Z. Down a group, the IE₁ value generally decreases with increasing Z. There are some systematic deviations from this trend, however. Note that the ionization energy of boron (atomic number 5) is less than that of beryllium (atomic number 4) even though the nuclear charge of boron is greater by one proton. This can be explained because the energy of the subshells increases as l increases, due to penetration and shielding (as discussed previously in this chapter). Within any one shell, the s electrons are lower in energy than the p electrons. This means that an s electron is harder to remove from an atom than a p electron in the same shell. The electron removed during the ionization of beryllium ([He]2s²) is an s electron, whereas the electron removed during the ionization of boron ([He]2s²2p¹) is a p electron; this results in a lower first ionization energy for boron, even though its nuclear charge is greater by one proton. Thus, we see a small deviation from the predicted trend occurring each time a new subshell begins.

Another deviation occurs as orbitals become more than one-half filled. The first ionization energy for oxygen is slightly less than that for nitrogen, despite the trend in increasing IE₁ values across a period. Looking at the orbital diagram of oxygen, we can see that removing one electron will eliminate the electron–electron repulsion caused by pairing the electrons in the 2p orbital and will result in a half-filled orbital (which is energetically favorable). Analogous changes occur in succeeding periods (note the dip for sulfur after phosphorus in Figure (PageIndex{4}).

Removing an electron from a cation is more difficult than removing an electron from a neutral atom because of the greater electrostatic attraction to the cation. Likewise, removing an electron from a cation with a higher positive charge is more difficult than removing an electron from an ion with a lower charge. Thus, successive ionization energies for one element always increase. As seen in Table (PageIndex{2}), there is a large increase in the ionization energies (color change) for each element. This jump corresponds to removal of the core electrons, which are harder to remove than the valence electrons. For example, Sc and Ga both have three valence electrons, so the rapid increase in ionization energy occurs after the third ionization.

Example (PageIndex{2}): Ranking Ionization Energies

Predict the order of increasing energy for the following processes: IE₁ for Al, IE₁ for Tl, IE₂ for Na, IE₃ for Al.

Solution

Amount Of Valence Electrons In Fluorine

Removing the 6p¹ electron from Tl is easier than removing the 3p¹ electron from Al because the higher n orbital is farther from the nucleus, so IE₁(Tl) < IE₁(Al). Ionizing the third electron from

[ce{Al}hspace{20px}ce{(Al^2+⟶Al^3+ + e- )} nonumber]

requires more energy because the cation Al²⁺ exerts a stronger pull on the electron than the neutral Al atom, so IE₁(Al) < IE₃(Al). The second ionization energy for sodium removes a core electron, which is a much higher energy process than removing valence electrons. Putting this all together, we obtain:

IE₁(Tl) < IE₁(Al) < IE₃(Al) < IE₂(Na).

Exercise (PageIndex{2})

Which has the lowest value for IE₁: O, Po, Pb, or Ba?

Ba

Variation in Electron Affinities

The electron affinity [EA] is the energy change for the process of adding an electron to a gaseous atom to form an anion (negative ion).

This process can be either endothermic or exothermic, depending on the element. The EA of some of the elements is given in Figure (PageIndex{6}). You can see that many of these elements have negative values of EA, which means that energy is released when the gaseous atom accepts an electron. However, for some elements, energy is required for the atom to become negatively charged and the value of their EA is positive. Just as with ionization energy, subsequent EA values are associated with forming ions with more charge. The second EA is the energy associated with adding an electron to an anion to form a –2 ion, and so on.

As we might predict, it becomes easier to add an electron across a series of atoms as the effective nuclear charge of the atoms increases. We find, as we go from left to right across a period, EAs tend to become more negative. The exceptions found among the elements of group 2 (2A), group 15 (5A), and group 18 (8A) can be understood based on the electronic structure of these groups. The noble gases, group 18 (8A), have a completely filled shell and the incoming electron must be added to a higher n level, which is more difficult to do. Group 2 (2A) has a filled ns subshell, and so the next electron added goes into the higher energy np, so, again, the observed EA value is not as the trend would predict. Finally, group 15 (5A) has a half-filled np subshell and the next electron must be paired with an existing np electron. In all of these cases, the initial relative stability of the electron configuration disrupts the trend in EA.

Number Of Valence Electrons In Chlorine

We also might expect the atom at the top of each group to have the largest EA; their first ionization potentials suggest that these atoms have the largest effective nuclear charges. However, as we move down a group, we see that the second element in the group most often has the greatest EA. The reduction of the EA of the first member can be attributed to the small size of the n = 2 shell and the resulting large electron–electron repulsions. For example, chlorine, with an EA value of –348 kJ/mol, has the highest value of any element in the periodic table. The EA of fluorine is –322 kJ/mol. When we add an electron to a fluorine atom to form a fluoride anion (F^–), we add an electron to the n = 2 shell. The electron is attracted to the nucleus, but there is also significant repulsion from the other electrons already present in this small valence shell. The chlorine atom has the same electron configuration in the valence shell, but because the entering electron is going into the n = 3 shell, it occupies a considerably larger region of space and the electron–electron repulsions are reduced. The entering electron does not experience as much repulsion and the chlorine atom accepts an additional electron more readily.

The properties discussed in this section (size of atoms and ions, effective nuclear charge, ionization energies, and electron affinities) are central to understanding chemical reactivity. For example, because fluorine has an energetically favorable EA and a large energy barrier to ionization (IE), it is much easier to form fluorine anions than cations. Metallic properties including conductivity and malleability (the ability to be formed into sheets) depend on having electrons that can be removed easily. Thus, metallic character increases as we move down a group and decreases across a period in the same trend observed for atomic size because it is easier to remove an electron that is farther away from the nucleus.

Summary

Electron configurations allow us to understand many periodic trends. Covalent radius increases as we move down a group because the n level (orbital size) increases. Covalent radius mostly decreases as we move left to right across a period because the effective nuclear charge experienced by the electrons increases, and the electrons are pulled in tighter to the nucleus. Anionic radii are larger than the parent atom, while cationic radii are smaller, because the number of valence electrons has changed while the nuclear charge has remained constant. Ionization energy (the energy associated with forming a cation) decreases down a group and mostly increases across a period because it is easier to remove an electron from a larger, higher energy orbital. Electron affinity (the energy associated with forming an anion) is more favorable (exothermic) when electrons are placed into lower energy orbitals, closer to the nucleus. Therefore, electron affinity becomes increasingly negative as we move left to right across the periodic table and decreases as we move down a group. For both IE and electron affinity data, there are exceptions to the trends when dealing with completely filled or half-filled subshells.

Number Of Valence Electrons In Fluorine

Glossary

Number Of Valence Electrons In Fluoride

covalent radius

one-half the distance between the nuclei of two identical atoms when they are joined by a covalent bond

effective nuclear charge

charge that leads to the Coulomb force exerted by the nucleus on an electron, calculated as the nuclear charge minus shielding

electron affinity

energy required to add an electron to a gaseous atom to form an anion

ionization energy

energy required to remove an electron from a gaseous atom or ion. The associated number (e.g., second ionization energy) corresponds to the charge of the ion produced (X²⁺)

Determine The Number Of Valence Electrons In Fluorine Atom

isoelectronic

group of ions or atoms that have identical electron configurations

Contributors and Attributions

• Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors. Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/85abf193-2bd...a7ac8df6@9.110).