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• How Many Neutrons Are In Fluorine
• Fluorine Number Of Electrons
• Fluorine Number Of Electrons In Valence Shell
• Fluorine Number Of Electrons Gained Or Lost
• Fluorine Number Of Electrons To Gain
• Number Of Protons Fluorine

• Fluorine has 7 valence electrons. What type of bond is likely to form between two atoms of fluorine? Bohr models show the number of electrons (regular and valence) which helps determine which sort of bond should take place. How many electrons can the first shell in an atom hold?
• Fluorine has an atomic number of 9, so it has 9 electrons. Become a more stable element it gains an electron to have a full This gives fluorine a charge of 1.
• PF3 is a tetra-atomic molecule where phosphorus donates three valence electrons, and three fluorine atoms accept one electron each to undergo a bond formation and reach a stable condition. Below are the steps to draw the lewis structure of the PF3 molecule 1. Find out the total number of valence electrons in PF3, which is 26.

Fluorine has the atomic number of 9. This means it has 9 protons in the nuclei of its atoms. So, a neutral fluorine atom would also have 9 electrons.

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Fluorine (F), most reactive chemical element and the lightest member of the halogen elements, or Group 17 (Group VIIa) of the periodic table. Its chemical activity can be attributed to its extreme ability to attract electrons (it is the most electronegative element) and to the small size of its atoms.

History

The fluorine-containing mineral fluorspar (or fluorite) was described in 1529 by the German physician and mineralogist Georgius Agricola. It appears likely that crude hydrofluoric acid was first prepared by an unknown English glassworker in 1720. In 1771 the Swedish chemist Carl Wilhelm Scheele obtained hydrofluoric acid in an impure state by heating fluorspar with concentrated sulfuric acid in a glass retort, which was greatly corroded by the product; as a result, vessels made of metal were used in subsequent experiments with the substance. The nearly anhydrous acid was prepared in 1809, and two years later the French physicist André-Marie Ampère suggested that it was a compound of hydrogen with an unknown element, analogous to chlorine, for which he suggested the name fluorine. Fluorspar was then recognized to be calcium fluoride.

The isolation of fluorine was for a long time one of the chief unsolved problems in inorganic chemistry, and it was not until 1886 that the French chemist Henri Moissan prepared the element by electrolyzing a solution of potassium hydrogen fluoride in hydrogen fluoride. He received the 1906 Nobel Prize for Chemistry for isolating fluorine. The difficulty in handling the element and its toxic properties contributed to the slow progress in fluorine chemistry. Indeed, up to the time of World War II the element appeared to be a laboratory curiosity. Then, however, the use of uranium hexafluoride in the separation of uranium isotopes, along with the development of organic fluorine compounds of industrial importance, made fluorine an industrial chemical of considerable use.

Occurrence and distribution

The fluorine-containing mineral fluorspar (fluorite, CaF₂) has been used for centuries as a flux (cleansing agent) in various metallurgical processes. The name fluorspar is derived from the Latin fluere, “to flow.” The mineral subsequently proved to be a source of the element, which was accordingly named fluorine. The colourless, transparent crystals of fluorspar exhibit a bluish tinge when illuminated, and this property is accordingly known as fluorescence.

Fluorine is found in nature only in the form of its chemical compounds, except for trace amounts of the free element in fluorspar that has been subjected to radiation from radium. Not a rare element, it makes up about 0.065 percent of Earth’s crust. The principal fluorine-containing minerals are (1) fluorspar, deposits of which occur in Illinois, Kentucky, Derbyshire, southern Germany, the south of France, and Russia and the chief source of fluorine, (2) cryolite (Na₃AlF₆), chiefly from Greenland, (3) fluoroapatite (Ca₅[PO₄]₃[F,Cl]), widely distributed and containing variable amounts of fluorine and chlorine, (4) topaz (Al₂SiO₄[F,OH]₂), the gemstone, and (5) lepidolite, a mica as well as a component of animal bones and teeth.

Physical and chemical properties

At room temperature fluorine is a faintly yellow gas with an irritating odour. Inhalation of the gas is dangerous. Upon cooling fluorine becomes a yellow liquid. There is only one stable isotope of the element, fluorine-19.

Because fluorine is the most electronegative of the elements, atomic groupings rich in fluorine are often negatively charged. Methyl iodide (CH₃I) and trifluoroiodomethane (CF₃I) have different charge distributions as shown in the following formulas, in which the Greek symbol δ indicates a partial charge:

The first ionization energy of fluorine is very high (402 kilocalories per mole), giving a standard heat formation for the F⁺ cation of 420 kilocalories per mole.

The small size of the fluorine atom makes it possible to pack a relatively large number of fluorine atoms or ions around a given coordination centre (central atom) where it forms many stable complexes—for example, hexafluorosilicate (SiF₆)²⁻ and hexafluoroaluminate (AlF₆)³⁻. Fluorine is the most powerfully oxidizing element. No other substance, therefore, is able to oxidize the fluoride anion to the free element, and for this reason the element is not found in the free state in nature. For more than 150 years, all chemical methods had failed to produce the element, success having been achieved only by the use of electrolytic methods. However, in 1986 American chemist Karl O. Christe reported the first chemical preparation of fluorine, where “chemical preparation” means a method that does not use techniques such as electrolysis, photolysis, and discharge or use fluorine itself in the synthesis of any of the starting materials. He used K₂MnF₆ and antimony pentafluoride (SbF₅), both of which can be easily prepared from HF solutions.

The high oxidizing power of fluorine allows the element to produce the highest oxidation numbers possible in other elements, and many high oxidation state fluorides of elements are known for which there are no other corresponding halides—e.g., silver difluoride (AgF₂), cobalt trifluoride (CoF₃), rhenium heptafluoride (ReF₇), bromine pentafluoride (BrF₅), and iodine heptafluoride (IF₇).

Fluorine (F₂), composed of two fluorine atoms, combines with all other elements except helium and neon to form ionic or covalent fluorides. Some metals, such as nickel, are quickly covered by a fluoride layer, which prevents further attack of the metal by the element. Certain dry metals, such as mild steel, copper, aluminum, or Monel (a 66 percent nickel, 31.5 percent copper alloy), are not attacked by fluorine at ordinary temperatures. For work with fluorine at temperatures up to 600 °C (1,100 °F), Monel is suitable; sintered alumina is resistant up to 700 °C (1,300 °F). When lubricants are required, fluorocarbon oils are most suitable. Fluorine reacts violently with organic matter (such as rubber, wood, and cloth), and controlled fluorination of organic compounds by the action of elemental fluorine is only possible if special precautions are taken.

Learning Objectives

• To understand the basics of electron shielding and penetration

For an atom or an ion with only a single electron, we can calculate the potential energy by considering only the electrostatic attraction between the positively charged nucleus and the negatively charged electron. When more than one electron is present, however, the total energy of the atom or the ion depends not only on attractive electron-nucleus interactions but also on repulsive electron-electron interactions. When there are two electrons, the repulsive interactions depend on the positions of both electrons at a given instant, but because we cannot specify the exact positions of the electrons, it is impossible to exactly calculate the repulsive interactions. Consequently, we must use approximate methods to deal with the effect of electron-electron repulsions on orbital energies. These effects are the underlying basis for the periodic trends in elemental properties that we will explore in this chapter.

Electron Shielding and Effective Nuclear Charge

If an electron is far from the nucleus (i.e., if the distance (r) between the nucleus and the electron is large), then at any given moment, many of the other electrons will be between that electron and the nucleus (Figure (PageIndex{1})). Hence the electrons will cancel a portion of the positive charge of the nucleus and thereby decrease the attractive interaction between it and the electron farther away. As a result, the electron farther away experiences an effective nuclear charge ((Z_{eff})) that is less than the actual nuclear charge (Z). This effect is called electron shielding.

As the distance between an electron and the nucleus approaches infinity, (Z_{eff}) approaches a value of 1 because all the other ((Z − 1)) electrons in the neutral atom are, on the average, between it and the nucleus. If, on the other hand, an electron is very close to the nucleus, then at any given moment most of the other electrons are farther from the nucleus and do not shield the nuclear charge. At (r ≈ 0), the positive charge experienced by an electron is approximately the full nuclear charge, or (Z_{eff} ≈ Z). At intermediate values of (r), the effective nuclear charge is somewhere between 1 and (Z):

[1 ≤ Z_{eff} ≤ Z.]

Notice that (Z_{eff} = Z) only for hydrogen and only for helium are (Z_{eff}) and (Z) comparable in magnitude (Figure (PageIndex{2})).

Shielding

Shielding refers to the core electrons repelling the outer electrons, which lowers the effective charge of the nucleus on the outer electrons. Hence, the nucleus has 'less grip' on the outer electrons insofar as it is shielded from them.

(Z_{eff}) can be calculated by subtracting the magnitude of shielding from the total nuclear charge and the effective nuclear charge of an atom is given by the equation:

[ Z_{eff}=Z-S label{4}]

where (Z) is the atomic number (number of protons in nucleus) and (S) is the shielding constant. The value of (Z_{eff}) will provide information on how much of a charge an electron actually experiences.

We can see from Equation ref{4} that the effective nuclear charge of an atom increases as the number of protons in an atom increases (Figure (PageIndex{2})). Therefore as we go from left to right on the periodic table the effective nuclear charge of an atom increases in strength and holds the outer electrons closer and tighter to the nucleus. As we will discuss later on in the chapter, this phenomenon can explain the decrease in atomic radii we see as we go across the periodic table as electrons are held closer to the nucleus due to increase in number of protons and increase in effective nuclear charge.

The shielding constant can be estimated by totaling the screening by all electrons ((n)) except the one in question.

[ S = sum_{i}^{n-1} S_i label{2.6.0}]

where (S_i) is the shielding of the i^th electron.

Electrons that are shielded from the full charge of the nucleus experience an effective nuclear charge ((Z_{eff})) of the nucleus, which is some degree less than the full nuclear charge an electron would feel in a hydrogen atom or hydrogenlike ion.

From Equations ref{4} and ref{2.6.0}, (Z_{eff}) for a specific electron can be estimated is the shielding constants for that electron of all other electrons in species is known. A simple approximation is that all other electrons shield equally and fully:

[S_i=1 label{simple}]

This crude approximation is demonstrated in Example (PageIndex{1}).

Example (PageIndex{1}): Fluorine, Neon, and Sodium

What is the effective attraction (Z_{eff}) experienced by the valence electrons in the three isoelectronic species: the fluorine anion, the neutral neon atom, and sodium cation?

Solution

Each species has 10 electrons, and the number of nonvalence electrons is 2 (10 total electrons - 8 valence), but the effective nuclear charge varies because each has a different atomic number (A). This is an application of Equations ref{4} and ref{2.6.0}. We use the simple assumption that all electrons shield equally and fully the valence electrons (Equation ref{simple}).

The charge (Z) of the nucleus of a fluorine atom is 9, but the valence electrons are screened appreciably by the core electrons (four electrons from the 1s and 2s orbitals) and partially by the 7 electrons in the 2p orbitals.

• (Z_mathrm{eff}(mathrm{F}^-) = 9 - 2 = 7+)
• (Z_mathrm{eff}(mathrm{Ne}) = 10 - 2 = 8+)
• (Z_mathrm{eff}(mathrm{Na}^+) = 11 - 2 = 9+)

So the sodium cation has the greatest effective nuclear charge. This also suggests that (mathrm{Na}^+) has the smallest radius of these species and that is correct.

Exercise (PageIndex{1}): Sodium Species

What is the effective attraction (Z_{eff}) experienced by the valence electrons in the sodium anion, the neutral sodium atom, and sodium cation? Use the simple approximation for shielding constants. Compare your result for the sodium atom to the more accurate value in Table (PageIndex{1}) and proposed an origin for the difference.

How Many Neutrons Are In Fluorine

• (Z_mathrm{eff}(mathrm{Na}^-) = 11 - 2 = 7+)
• (Z_mathrm{eff}(mathrm{Na}) = 11 - 2 = 8+)
• (Z_mathrm{eff}(mathrm{Na}^+) = 11 - 2 = 9+)

The (Z_{eff}) in Table (PageIndex{1}) for (Z_mathrm{eff}(mathrm{Na}) is 10.63 and appreciables larger than the 8 estimated above. This means the simple approximation (Equation ref{simple}) overestimates the shielding constant (S).

Electron Penetration

The approximation in Equation ref{simple} is a good first order description of electron shielding, but the actual (Z_{eff}) experienced by an electron in a given orbital depends not only on the spatial distribution of the electron in that orbital but also on the distribution of all the other electrons present. This leads to large differences in (Z_{eff}) for different elements, as shown in Figure (PageIndex{2}) for the elements of the first three rows of the periodic table.

Penetration describes the proximity to which an electron can approach to the nucleus. In a multi-electron system, electron penetration is defined by an electron's relative electron density (probability density) near the nucleus of an atom (Figure (PageIndex{3})). Electrons in different orbitals have different electron densities around the nucleus. In other words, penetration depends on the shell ((n)) and subshell ((l)).

For example, a 1s electron (Figure (PageIndex{3}); purple curve) has greater electron density near the nucleus than a 2p electron (Figure (PageIndex{3}); red curve) and has a greater penetration. This related to the shielding constants since the 1s electrons are closer to the nucleus than a 2p electron, hence the 1s screens a 2p electron almost perfectly ((S=1). However, the 2s electron has a lower shielding constant ((S<1) because it can penetrate close to the nucleus in the small area of electron density within the first spherical node (Figure (PageIndex{3}); green curve). In this way the 2s electron can 'avoid' some of the shielding effect of the inner 1s electron.

For the same shell value ((n)) the penetrating power of an electron follows this trend in subshells (Figure (PageIndex{3})):

Fluorine Number Of Electrons

[s > p > d approx f. label{better1}]

for different values of shell (n) and subshell (l), penetrating power of an electron follows this trend:

[ce{1s > 2s > 2p > 3s > 3p > 4s > 3d > 4p > 5s > 4d > 5p > 6s > 4f ...} label{better2}]

Penetration

Penetration describes the proximity of electrons in an orbital to the nucleus. Electrons that have greater penetration can get closer to the nucleus and effectively block out the charge from electrons that have less proximity.

Data from E. Clementi and D. L. Raimondi; The Journal of Chemical Physics 38, 2686 (1963).

Fluorine Number Of Electrons In Valence Shell

Because of the effects of shielding and the different radial distributions of orbitals with the same value of n but different values of l, the different subshells are not degenerate in a multielectron atom. For a given value of n, the ns orbital is always lower in energy than the np orbitals, which are lower in energy than the nd orbitals, and so forth. As a result, some subshells with higher principal quantum numbers are actually lower in energy than subshells with a lower value of n; for example, the 4s orbital is lower in energy than the 3d orbitals for most atoms.

A Better Estimation of Shielding: Slater Rules

Fluorine Number Of Electrons Gained Or Lost

The concepts of electron shielding, orbital penetration and effective nuclear charge were introduced above, but we did so in a qualitative manner (e.g., Equations ref{better1} and ref{better2}). A more accurate model for estimating electron shielding and corresponding effective nuclear charge experienced is Slater's Rules. However, the application of these rules is outside the scope of this text.

Fluorine Number Of Electrons To Gain

Summary

Number Of Protons Fluorine

The calculation of orbital energies in atoms or ions with more than one electron (multielectron atoms or ions) is complicated by repulsive interactions between the electrons. The concept of electron shielding, in which intervening electrons act to reduce the positive nuclear charge experienced by an electron, allows the use of hydrogen-like orbitals and an effective nuclear charge ((Z_{eff})) to describe electron distributions in more complex atoms or ions. The degree to which orbitals with different values of l and the same value of n overlap or penetrate filled inner shells results in slightly different energies for different subshells in the same principal shell in most atoms.