The number of electrons in an electrically-neutral atom is the same as the number of protons in the nucleus. Therefore, the number of electrons in neutral atom of Fluorine is 9. Each electron is influenced by the electric fields produced by the positive nuclear charge and the other (Z – 1) negative electrons in the atom. Fluorine (F) has an atomic number of 9 and by referring to a periodic table, we can see that fluorine (F) is in Group 7A. Elements in Group 7A all have 7 valence electrons. To illustrate this, we.

Chemical Bonding

Why do chemical bonds form?In large part, it is to lower the potential energy (PE) of the system.Potential energy arises from the interaction of positive and negative charges.At an atomic level, positive charges are carried by protons and negative charges are carried by electrons. The PE can be calculated using Coulomb's Law, which is the product of two charges, Q₁ and Q₂ divided by the distance between the charges, d.If the two charges have the same sign (+,+ or -,-) the PE will be a positive number.Like charges repel each other, so positive PE is a destabilizing factor.If the two charges have different signs, the PE will be negative.This indicates an attractive force between the charges and is a stabilizing factor.Chemical bonding leads to a lowering of the PE and formation of more stable chemical species.

Ionic bonding

Ionic bonds form between metals and non-metals.Metals are the elements on the left side of the Periodic Table.The most metallic elements are Cesium and Francium.Metals tend to lose electrons to attain Noble Gas electron configuration.Groups 1 and 2 (the active metals) lose 1 and 2 valence electrons, respectively, because of their low Ionization energies.

Non-metals are limited to the elements in the upper right hand corner of the Periodic Table.The most non-metallic element is fluorine.Non-metals tend to gain electrons to attain Noble Gas configurations.The have relatively high Electron affinities and high Ionization energies.

Metals tend to lose electrons and non-metals tend to gain electrons, so in reactions involving these two groups, there is electron transfer from the metal to the non-metal.The metal is oxidized and the non-metal is reduced.An example of this is the reaction between the metal, sodium, and the non-metal, chlorine.The sodium atom gives up an electron to form the Na⁺ ion and the chorine molecule gains electrons to form 2 Cl^- ions. The charges on these anions and cations are stabilized by forming a crystal lattice, in which each of the ions is surrounded by counter ions.

The sodium ions, Na⁺, are represented by the red spheres, and the chloride ions, Cl^-, by the yellow spheres.The formula for the product, NaCl, indicates the ratio of sodium ions to chloride ions.There are no individual molecules of NaCl.

Covalent Bonding

Covalent bonding takes place between non-metals.There is no transfer of electrons, but a sharing of valence electrons.The non-metals all have fairly high ionization energies, meaning that it is relatively difficult to remove their valence electrons.The non-metals also have relatively high electron affinities, so they tend to attract electrons to themselves.So, they share valence electrons with other non-metals.The shared electrons are held between the two nuclei.The formula of covalent compounds represents actual numbers of atoms that are bonded to form molecules, like C₆H₁₂O₆ for glucose.Covalent species exist as individual molecules.

Metallic Bonding

Metallic bonding exists between metal atoms.Metals have relatively low ionization energies (easily removed electrons) but also low electron affinities (very little tendency to gain electrons).So, metals will share electrons.However, it is a different sort of bonding than covalent bonding.Metals share valence electrons, but these are not localized between individual atoms.Instead, they are distributed throughout the metal and are completely delocalized.They are often described as being a 'sea' of electrons which flow freely between the atoms.The graphic, below, attempts to show this.The darker gray spheres are the metal nuclei and core electrons.The lighter gray areas are the loosely held valence electrons, which are effectively shared by all Extended serial monitor. of the metal atoms.

Ionic bonding - Lattice Energy

Metals and non-metals interact to form ionic compounds.An example of this is the reaction between Na and Cl₂.

2 Na (s) + Cl₂ (g) → 2 NaCl (s)

The link, below (which sometimes works and sometimes doesn't) shows this reaction taking place.

It is an extremely exothermic reaction.A great deal of heat is given off, indicating a large decrease in the PE of the system.The product, NaCl, is much more stable than the reactants, Na and Cl₂.We can break this reaction down into a few steps, to try to figure out there all of this energy comes from.We expect a large negative number as the final answer. Things 3 mac free download.

First, we need to ionize sodium and chlorine:

Ionization Energy Na → Na⁺ + e^-+496 kJ/molEnergy needs to be added in order to remove the electron

Electron AffinityCl + e^- → Cl^--349 kJ/molEnergy is given off when chlorine gains an electron

The sum of these two is positive.There must be some other step involved.

That step involves assembling the ions into a crystal lattice, so it is called the Lattice Energy.

Lattice energy = E = k Q₁ Q₂

d

The Q values are ionic charges, k is a constant that depends on the units used and d is the dielectric of the medium. For NaCl, this equals -504 kJ/mol.So, the overall energy change is negative.

Lewis electron-dot symbols

Lewis electron-dot symbols represent the valence electrons on each atom.The element symbol itself, represents the nucleus and core electrons and each 'dot' represents a valence electron. Notice that Hund's rule is obeyed, and electron are kept unpaired if possible.

With the metals, (to the left of the red line) the total number of dots represents an electron that the element can lose in order to form a cation.In the non-metals (to the right of the red line) the number of unpaired dot represents the number of electrons that can become paired, through the gain or sharing of electrons.So, the number of unpaired dots equals either the negative charge on the anion that forms, from electron transfer with a metal, or the number of covalent bonds that the element can form by sharing electrons with other non-metals.So, Mg, with two dots, tends to form the Mg²⁺ ion.Carbon, with 4 unpaired dots, can form the carbide ion, C^4-, when reacting with metals, or can form four bonds when reacting with non-metals.

The reaction between Na and Cl₂ can be written in terms of their Lewis electron dot structures.

2 Na (s) + Cl₂ (g) → 2 NaCl (s)

Chlorine gains one valence electron to form Cl^- and sodium loses one electron to form Na⁺.Both now have Noble gas electron configurations.

Ionic radii

When atoms lose electrons to form cations, the ionic radius is always smaller than the atomic radius.There are fewer electrons, with an unchanged nuclear charge, Z.This means that the remaining electrons will be held more strongly and more closely to the nucleus.When atoms gain electrons to forms anions, the ionic radius is always larger than the atomic radius.Shown below is a chart of ionic radii.

Elemental sodium is larger than elemental chlorine.However, when they are ionized, their relative sizes reverse.It is very difficult to predict absolute sizes.Relative sizes can be predicted for isoelectronic series, species which have the same number of electrons.For example O^2- and F^- both have 10 electrons.The nuclear charge on oxygen is +8 and the nuclear charge on fluorine is +9.The positive charges increase, but the negative charges stay the same (-10).So, F^- will be smaller due to the increased attraction (+9/-10 versus +8/-10).The series of In³⁺, Sn⁴⁺ and Sb⁵⁺ show the same trend.They all have 46e^-, but have nuclear charges of +49, +50 and +51, respectively.Sb⁵⁺ is the smallest of the three.

Covalent bonds

In covalent bonds, valence electrons are shared between bonded atoms. For example, the element hydrogen, exists as H₂ gas. Its structure can be described using the Lewis dot structure for hydrogen:

Each hydrogen atom now has a share in two valence electrons, giving it the electron configuration of helium. A pair of shared electrons is a chemical bond. Hydrogen is usually represented as H-H, where the line represents a bonding pair of electrons.

The element fluorine exists as the diatomic gas, F₂. Again, the Lewis structure illustrates the bonding.

This would be represented as F-F. The non-bonding electrons (6 on each F) are called lone-pair electrons.

Covalent bonds also forms between different atoms, as in hydrofluoric acid;.

The hydrogen atom is bonded to the fluorine atom with bonding pairs of electrons, and the fluorine also has three lone pairs of electrons.

In this case, the electrons are not shared equally between the fluorine and the hydrogen, but spend more time near the fluorine atom, giving it a partial negative charge (δ^-) and putting a partial positive charge (δ⁺) on the hydrogen. There is more electron density in the red regions, than in the blue regions.

This type of bond is called a polar covalent bond. It is intermediate between an ionic bond, in which the transfer of the electron is complete and a purely covalent bond (like F₂ and H₂) in which the sharing of the bonding electrons is exactly equal.

Atoms of the same element can differ in. The extent to which electrons will be drawn toward an atom in a molecule is determined by the electronegativity values of the bonded atoms.

Electronegativity

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.

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

Valence Electron Configuration Of Fluorine

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)

Total Number Of Valence Electrons In Fluorine

( 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.

The Number Of Valence Electrons In Fluorine

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

Total Number Of Valence Electrons In Fluorine

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

Total Number Of Valence Electrons In Fluorine Atom

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.