Ionization Energy Of Fluorine

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The unity for ionization energy is eV. Please note that the elements do not show their natural relation towards each other as in the Periodic system. There you can find the metals, semi-conductor(s), non-metal(s), inert noble gas(ses), Halogens, Lanthanoides, Actinoids (rare earth elements) and transition metals. Ionization energy is the energy needed to take away an electron from an atom; the trend = it is the lowest in the bottom left and increases toward the upper right Why does fluorine have a higher ionization energy than iodine? A Fluorine atom, for example, requires the following ionization energy to remove the outermost electron. F + IE → F + + e − IE = 17.4228 eV. The ionization energy associated with removal of the first electron is most commonly used. The nth ionization energy refers to the amount of energy required to remove an electron from the species with. However oxygen has greater second ionization energy than fluorine and also nitrogen. Reason: Since Oxygen atom gets stable electronic configuration, 2s 2 2p 3 after removing one electron, the O + shows greater ionization energy than F + as well as N +. Ionization energy determinations. IE (eV) Method., High Resolution Determination of the Electron Affinity of Fluorine and Bromine using Crossed Ion.


The halogens are located on the left of the noble gases on the periodic table. These five toxic, non-metallic elements make up Group 17 of the periodic table and consist of: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). Although astatine is radioactive and only has short-lived isotopes, it behaves similar to iodine and is often included in the halogen group. Because the halogen elements have seven valence electrons, they only require one additional electron to form a full octet. This characteristic makes them more reactive than other non-metal groups.


Halogens form diatomic molecules (of the form X2​, where X denotes a halogen atom) in their elemental states. The bonds in these diatomic molecules are non-polar covalent single bonds. However, halogens readily combine with most elements and are never seen uncombined in nature. As a general rule, fluorine is the most reactive halogen and astatine is the least reactive. All halogens form Group 1 salts with similar properties. In these compounds, halogens are present as halide anions with charge of -1 (e.g. Cl-, Br-, etc.). Replacing the -ine ending with an -ide ending indicates the presence of halide anions; for example, Cl- is named 'chloride.' In addition, halogens act as oxidizing agents—they exhibit the property to oxidize metals. Therefore, most of the chemical reactions that involve halogens are oxidation-reduction reactions in aqueous solution. The halogens often form single bonds, when in the -1 oxidation state, with carbon or nitrogen in organic compounds. When a halogen atom is substituted for a covalently-bonded hydrogen atom in an organic compound, the prefix halo- can be used in a general sense, or the prefixes fluoro-, chloro-, bromo-, or iodo- can be used for specific halogen substitutions. Halogen elements can cross-link to form diatomic molecules with polar covalent single bonds.

Chlorine (Cl2) was the first halogen to be discovered in 1774, followed by iodine (I2), bromine (Br2), fluorine (F2), and astatine (At, discovered last in 1940). The name 'halogen' is derived from the Greek roots hal- ('salt') and -gen ('to form'). Together these words combine to mean 'salt former', referencing the fact that halogens form salts when they react with metals. Halite is the mineral name for rock salt, a natural mineral consisting essentially of sodium chloride (NaCl). Lastly, the halogens are also relevant in daily life, whether it be the fluoride that goes in toothpaste, the chlorine that disinfects drinking water, or the iodine that facilitates the production of thyroid hormones in one's body.


Fluorine - Fluorine has an atomic number of 9 and is denoted by the symbol F. Elemental fluorine was first discovered in 1886 by isolating it from hydrofluoric acid. Fluorine exists as a diatomic molecule in its free state (F2) and is the most abundant halogen found in the Earth's crust. Fluorine is the most electronegative element in the periodic table. It appears as a pale yellow gas at room temperature. Fluorine also has a relatively small atomic radius. Its oxidation state is always -1 except in its elemental, diatomic state (in which its oxidation state is zero). Fluorine is extremely reactive and reacts directly with all elements except helium (He), neon (Ne) and argon (Ar). In H2O solution, hydrofluoric acid (HF) is a weak acid. Although fluorine is highly electronegative, its electronegativity does not determine its acidity; HF is a weak acid due to the fact that the fluoride ion is basic (pH>7). In addition, fluorine produces very powerful oxidants. For example, fluorine can react with the noble gas xenon and form the strong oxidizing agent Xenon Difluoride (XeF2). There are many uses for fluorine, which will be discussed in Part VI of this article.

Chlorine - Chlorine has the atomic number 17 and the chemical symbol Cl. Chlorine was discovered in 1774 by extracting it from hydrochloric acid. In its elemental state, it forms the diatomic molecule Cl2. Chlorine exhibits multiple oxidation states, such as -1, +1, 3, 5, and 7. At room temperature it appears as a light green gas. Since the bond that forms between the two chlorine atoms is weak, the Cl2 molecule is very reactive. Chlorine reacts with metals to produce salts called chlorides. Chloride ions are the most abundant ions that dissolve in the ocean. Chlorine also has two isotopes: 35Cl and 37Cl. Sodium chloride is the most prevalent compound of the chlorides.

Bromine - Bromine has an atomic number of 35 with a symbol of Br. It was first discovered in 1826. In its elemental form, it is the diatomic molecule Br2. At room temperature, bromine is a reddish- brown liquid. Its oxidation states vary from -1, +1, 3, 4 and 5. Bromine is more reactive than iodine, but not as reactive as chlorine. Also, bromine has two isotopes: 79Br and 81Br. Bromine consists of bromide salts, which have been found in the sea. The world production of bromide has increased significantly over the years, due to its access and longer existence. Like all of the other halogens, bromine is an oxidizing agent, and is very toxic.

Iodine - Iodine has the atomic number 53 and symbol I. Iodine has oxidation states -1, +1, 5 and 7. Iodine exists as a diatomic molecule, I2, in its elemental state. At room temperature, it appears as a violet solid. Iodine has one stable isotope: 127I. It was first discovered in 1811 through the use of seaweed and sulfuric acid. Currently, iodide ions can be isolated in seawater. Although iodine is not very soluble in water, the solubility may increase if particular iodides are mixed in the solution. Iodine has many important roles in life, including thyroid hormone production. This will be discussed in Part VI of the text.

Astatine - Astatine is a radioactive element with an atomic number of 85 and symbol At. Its possible oxidation states include: -1, +1, 3, 5 and 7. It is the only halogen that is not a diatomic molecule and it appears as a black, metallic solid at room temperature. Astatine is a very rare element, so there is not that much known about this element. In addition, astatine has a very short radioactive half-life, no longer than a couple of hours. It was discovered in 1940 by synthesis. Also, it is thought that astatine is similar to iodine. However, these two elements are assumed to differ by their metallic character.

Table 1.1: Electron configurations of the halogens.
HalogenElectronic Configuration
Fluorine1s2 2s2 2p5
Chlorine [Ne]3s2 3p5
Bromine [Ar]3d104s24p5
Iodine [Kr]4d10 5s2 5p5
Astatine[Xe]4f14 5d10 6s2 6p5

Periodic Trends

The periodic trends observed in the halogen group:

Melting and Boiling Points (increases down the group)

The melting and boiling points increase down the group because of the van der Waals forces. The size of the molecules increases down the group. This increase in size means an increase in the strength of the van der Waals forces.

[F < Cl < Br < I < At]

Table 1.2: Melting and Boiling Points of Halogens
HalogenMelting Point (˚C)Boiling Point (˚C)

Atomic Radius (increases down the group)

The size of the nucleus increases down a group (F < Cl < Br < I < At) because the numbers of protons and neutrons increase. In addition, more energy levels are added with each period. This results in a larger orbital, and therefore a longer atomic radius.

Table 1.3: Atomic Radii of Halogens
HalogenCovalent Radius (pm)Ionic (X-) radius (pm)

Ionization Energy (decreases down the group)

If the outer valence electrons are not near the nucleus, it does not take as much energy to remove them. Therefore, the energy required to pull off the outermost electron is not as high for the elements at the bottom of the group since there are more energy levels. Also, the high ionization energy makes the element appear non-metallic. Iodine and astatine display metallic properties, so ionization energy decreases down the group (At < I < Br < Cl < F).

Table 1.4 Ionization Energy of Halogens
HalogenFirst Ionization Energy (kJ/mol)

Electronegativity (decreases down the group)

The number of valence electrons in an atom increases down the group due to the increase in energy levels at progressively lower levels. The electrons are progressively further from the nucleus; therefore, the nucleus and the electrons are not as attracted to each other. An increase in shielding is observed. Electronegativity therefore decreases down the group (At < I < Br < Cl < F).

Table 1.5: Electronegativity of Halogens

Electron Affinity (decreases down the group)

Since the atomic size increases down the group, electron affinity generally decreases (At < I < Br < F < Cl). An electron will not be as attracted to the nucleus, resulting in a low electron affinity. However, fluorine has a lower electron affinity than chlorine. This can be explained by the small size of fluorine, compared to chlorine.

Table 1.6: Electron Affinity of Halogens
HalogenElectron Affinity (kJ/mol)

Reactivity of Elements (decreases down the group)

The reactivities of the halogens decrease down the group ( At < I < Br < Cl < F). This is due to the fact that atomic radius increases in size with an increase of electronic energy levels. This lessens the attraction for valence electrons of other atoms, decreasing reactivity. This decrease also occurs because electronegativity decreases down a group; therefore, there is less electron 'pulling.' In addition, there is a decrease in oxidizing ability down the group.

Hydrogen Halides and Halogen Oxoacids

Hydrogen Halides

A halide is formed when a halogen reacts with another, less electronegative element to form a binary compound. Hydrogen, for example, reacts with halogens to form halides of the form HX:

  • Hydrogen Fluoride: HF
  • Hydrogen Chloride: HCl
  • Hydrogen Bromide: HBr
  • Hydrogen Iodide: HI

Hydrogen halides readily dissolve in water to form hydrohalic (hydrofluoric, hydrochloric, hydrobromic, hydroiodic) acids. The properties of these acids are given below:

  • The acids are formed by the following reaction: HX (aq) + H2O (l) X- (aq) + H3O+ (aq)
  • All hydrogen halides form strong acids, except HF
  • The acidity of the hydrohalic acids increases as follows: HF < HCl < HBr < HI

Hydrofluoric acid can etch glass and certain inorganic fluorides over a long period of time.

Ionization Energy Of Chlorine Kj/mol

It may seem counterintuitive to say that HF is the weakest hydrohalic acid because fluorine has the highest electronegativity. However,​ the H-F bond is very strong; if the H-X bond is strong, the resulting acid is weak. A strong bond is determined by a short bond length and a large bond dissociation energy. Of all the hydrogen halides, HF has the shortest bond length and largest bond dissociation energy.

Halogen Oxoacids

A halogen oxoacid is an acid with hydrogen, oxygen, and halogen atoms. The acidity of an oxoacid can be determined through analysis of the compound's structure. The halogen oxoacids are given below:

  • Hypochlorous Acid: HOCl
  • Chlorous Acid: HClO2
  • Chloric Acid: HClO3
  • Perchloric Acid: HClO4
  • Hypobromous Acid: HOBr
  • Bromic Acid: HBrO3
  • Perbromic Acid: HBrO4
  • Hypoiodous Acid: HOI
  • Iodic Acid: HIO3
  • Metaperiodic Acid: HIO4; H5IO6

In each of these acids, the proton is bonded to an oxygen atom; therefore, comparing proton bond lengths is not useful in this case. Instead, electronegativity is the dominant factor in the oxoacid's acidity. Acidic strength increases with more oxygen atoms bound to the central atom.

States of Matter at Room Temperature

Table 1.7: States of Matter and Appearance of Halogens
States of Matter (at Room Temperature)Halogen Appearance
Solid IodineViolet
AstatineBlack/Metallic [Assumed]
Liquid BromineReddish-Brown
Gas FluorinePale Yellow-Brown
ChlorinePale Green

Explanation for Appearance

The halogens' colors are results of the absorption of visible light by the molecules, which causes electronic excitation. Fluorine absorbs violet light, and therefore appears light yellow. Iodine, on the other hand, absorbs yellow light and appears violet (yellow and violet are complementary colors, which can be determined using a color wheel). The colors of the halogens grow darker down the group:

  • Fluorine pale yellow/brown
  • Chlorine pale green
  • Bromine red-brown
  • Iodine violet
  • Astatine* black/metallic

In closed containers, liquid bromine and solid iodine are in equilibrium with their vapors, which can often be seen as colored gases. Although the color for astatine is unknown, it is assumed that astatine must be darker than iodine's violet (i.e. black) based on the preceding trend.

Oxidation States of Halogens in Compounds

As a general rule, halogens usually have an oxidation state of -1. However, if the halogen is bonded to oxygen or to another halogen, it can adopt different states: the -2 rule for oxygen takes precedence over this rule; in the case of two different halogens bonded together, the more electronegative atom takes precedence and adopts the -1 oxidation state.

Example 1.1: Iodine Chloride (ICl)

Chlorine has an oxidation state of -1, and iodine will have an oxidation of +1. Chlorine is more electronegative than iodine, therefore giving it the -1 oxidation state.


Oxygen has a total oxidation state of -8 (-2 charge x 4 atoms= -8 total charge). Hydrogen has a total oxidation state of +1. Adding both of these values together, the total oxidation state of the compound so far is -7. Since the final oxidation state of the compound must be 0, bromine's oxidation state is +7.

One third exception to the rule is this: if a halogen exists in its elemental form (X2), its oxidation state is zero.

Table 1.8: Oxidation States of Halogens
HalogenOxidation States in Compounds
Fluorine(always) -1*
Chlorine-1, +1, +3, +5, +7
Bromine-1, +1, +3, +4, +5
Iodine-1, +1,+5, +7
Astatine-1, +1, +3, +5, +7

Example 1.3: Fluorine

Why does fluorine always have an oxidation state of-1 in its compounds?


Electronegativity increases across a period, and decreases down a group. Therefore, fluorine has the highest electronegativity of all of the elements, indicated by its position on the periodic table. Its electron configuration is 1s​2 2s2 2p5. If fluorine gains one more electron, the outermost p orbitals are completely filled (resulting in a full octet). Because fluorine has a high electronegativity, it can easily remove the desired electron from a nearby atom. Fluorine is then isoelectronic with a noble gas (with eight valence electrons); all its outermost orbitals are filled. Fluorine is much more stable in this state.

Applications of Halogens

Fluorine: Although fluorine is very reactive, it serves many industrial purposes. For example, it is a key component of the plastic polytetrafluoroethylene (called Teflon-TFE by the DuPont company) and certain other polymers, often referred to as fluoropolymers. Chlorofluorocarbons (CFCs) are organic chemicals that were used as refrigerants and propellants in aerosols before growing concerns about their possible environmental impact led to their discontinued use. Hydrochlorofluorocarbons (HFCs) are now used instead. Fluoride is also added to toothpaste and drinking water to help reduce tooth decay. Fluorine also exists in the clay used in some ceramics. Fluorine is associated with generating nuclear power as well. In addition, it is used to produce fluoroquinolones, which are antibiotics. Below is a list of some of fluorine's important inorganic compounds.

Table 1.9: Important Inorganic Compounds of Fluorine
Na3AlF6Manufacture of aluminum
CaF2Optical components, manufacture of HF, metallurgical flux
ClF3Fluorinating agent, reprocessing nuclear fuels
HFManufacture of F2, AlF3, Na3AlF6, and fluorocarbons
LiFCeramics manufacture, welding, and soldering
NaFFluoridating water, dental prophylaxis, insecticide
SF6Insulating gas for high-voltage electrical equipment
SnF2Manufacture of toothpaste
UF6 Manufacture of uranium fuel for nuclear reactors

Chlorine: Chlorine has many industrial uses. It is used to disinfect drinking water and swimming pools. Sodium hypochlorite (NaClO) is the main component of bleach. Hydrochloric acid, sometimes called muriatic acid, is a commonly used acid in industry and laboratories. Chlorine is also present in polyvinyl chloride (PVC), and several other polymers. PVC is used in wire insulation, pipes, and electronics. In addition, chlorine is very useful in the pharmaceutical industry. Medicinal products containing chlorine are used to treat infections, allergies, and diabetes. The neutralized form of hydrochloride is a component of many medications. Chlorine is also used to sterilize hospital machinery and limit infection growth. In agriculture, chlorine is a component of many commercial pesticides: DDT (dichlorodiphenyltrichloroethane) was used as an agricultural insecticide, but its use was discontinued.

Bromine: Bromine is used in flame retardants because of its fire-resistant properties. It also found in the pesticide methyl bromide, which facilitates the storage of crops and eliminates the spread of bacteria. However, the excessive use of methyl bromide has been discontinued due to its impact on the ozone layer. Bromine is involved in gasoline production as well. Other uses of bromine include the production of photography film, the content in fire extinguishers, and drugs treating pneumonia and Alzheimer's disease.

Iodine: Iodine is important in the proper functioning of the thyroid gland of the body. If the body does not receive adequate iodine, a goiter (enlarged thyroid gland) will form. Table salt now contains iodine to help promote proper functioning of the thyroid hormones. Iodine is also used as an antiseptic. Solutions used to clean open wounds likely contain iodine, and it is commonly found in disinfectant sprays. In addition, silver iodide is important for photography development.

Astatine: Because astatine is radioactive and rare, there are no proven uses for this halogen element. However, there is speculation that this element could aid iodine in regulating the thyroid hormones. Also, 211At has been used in mice to aid the study of cancer.

VII. Outside Links

  • Grube, Karl; Leffler, Amos J. 'Synthesis of metal halides (ML).' J. Chem. Educ.1993, 70, A204.
  • This video provides information about some of the physical properties of chlorine, bromine, and iodine:
  • The following video compares four halogens: fluorine, chlorine, bromine and iodine in terms of chemical reactions and physical properties.
  • Color wheel referenced to in the text:
  • Elson, Jesse. 'A bonding parameter. III, Water solubilities and melting points of the alkali halogens.' J. Chem. Educ.1969, 46, 86.
  • Fessenden, Elizabeth. 'Structural chemistry of the interhalogen compounds.' J. Chem. Educ. 1951, 28, 619.
  • Holbrook, Jack B.; Sabry-Grant, Ralph; Smith, Barry C.; Tandel, Thakor V. 'Lattice enthalpies of ionic halides, hydrides, oxides, and sulfides: Second-electron affinities of atomic oxygen and sulfur.' J. Chem. Educ. 1990, 67, 304.
  • Kildahl, Nicholas K. 'A procedure for determining formulas for the simple p-block oxoacids.' J. Chem. Educ. 1991, 68, 1001.
  • Liprandi, Domingo A.; Reinheimer, Orlando R.; Paredes, José F.; L'Argentière, Pablo C. 'A Simple, Safe Way To Prepare Halogens and Study Their Visual Properties at a Technical Secondary School.' J. Chem. Educ. 1999 76.
  • Meek, Terry L. 'Acidities of oxoacids: Correlation with charge distribution.'J. Chem. Educ. 1992, 69, 270.

Practice Problems

  1. Why does fluorine always have an oxidation state of -1 in its compounds?
  2. Find the oxidation state of the halogen in each problem:
    1. HOCl
    2. KIO3
    3. F2
  3. What are three uses of chlorine?
  4. Which element(s) exist(s) as a solid in room temperature?
  5. Do the following increase or decrease down the group of halogens?
    1. boiling point and melting point
    2. electronegativity
    3. ionization energy


  1. Electronegativity increases across a period, and decreases down a group. Therefore, fluorine has the highest electronegativity out of all of the elements. Because fluorine has seven valence electrons, it only needs one more electron to acheive a noble gas configuration (eight valence electrons). Therefore, it will be more likely to pull off an electron from a nearby atom.
  2. disinfecting water, pesticides, and medicinal products
    1. +1 (Hydrogen has an oxidation state of +1, and oxygen has an oxidation state of -2. Therefore, chlorine must have an oxidation state of +1 so that the total charge can be zero)
    2. +5 (Potassium's oxidation state is +1. Oxygen has an oxidation state of -2, so for this compound it is -6 (-2 charge x 3 atoms= -6). Since the total oxidation state has to be zero, iodine's oxidation state must be +5).
    3. 0 (Elemental forms always have an oxidation state of 0.)
  3. iodine and astatine
    1. increases
    2. decreases
    3. decreases


  1. Hill, Graham, and John Holman. Chemistry in Context. 5th ed. United Kingdom: Nelson Thornes, 2000. 224-25.
  2. Petrucci, Ralph H. Genereal Chemistry: Principles and Modern Applications. 9th Ed. New Jersey: Pearson Education Inc, 2007. 920-928.
  3. Verma, N.K., B. Kapila, and S.K. Khanna. Comprehensive Chemistry XII. New Delhi: Laxmi Publications, 2007. 718-30.

Ionization Energy and Electron Affinity

The First Ionization EnergyPatterns In First Ionization Energies
Exceptions to the General Pattern of First Ionization Energies2nd, 3rd, 4th, and Higher Ionization Energies
Electron AffinityConsequences of the Relative Size of Ionization Energies and Electron Affinities

The energy needed to remove one or more electrons from a neutral atom to form a positively charged ion is a physical property that influences the chemical behavior of the atom. By definition, the first ionization energy of an element is the energy needed to remove the outermost, or highest energy, electron from a neutral atom in the gas phase.

The process by which the first ionization energy of hydrogen is measured would be represented by the following equation.

Practice Problem 3:

Use the Bohr model to calculate the wavelength and energy of the photon that would have to be absorbed to ionize a neutral hydrogen atom in the gas phase.

The magnitude of the first ionization energy of hydrogen can be brought into perspective by comparing it with the energy given off in a chemical reaction. When we burn natural gas, about 800 kJ of energy is released per mole of methane consumed.

CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g)Ho = -802.4 kJ/mol

The thermite reaction, which is used to weld iron rails, gives off about 850 kJ of energy per mole of iron oxide consumed.

Fe2O3(s) + 2 Al(s) Al2O3(s) + 2 Fe(l)Ho = -851.5 kJ/mol

How To Calculate Ionization Energy

The first ionization energy of hydrogen is half again as large as the energy given off in either of these reactions.

The first ionization energy for helium is slightly less than twice the ionization energy for hydrogen because each electron in helium feels the attractive force of two protons, instead of one.

It takes far less energy, however, to remove an electron from a lithium atom, which has three protons in its nucleus.

Li(g) Li+(g) + e-Ho = 572.3 kJ/mol

This can be explained by noting that the outermost, or highest energy, electron on a lithium atom is in the 2s orbital. Because the electron in a 2s orbital is already at a higher energy than the electrons in a 1s orbital, it takes less energy to remove this electron from the atom.

The first ionization energies for the main group elements are given in the two figures below.

Two trends are apparent from these data.

  • In general, the first ionization energy increases as we go from left to right across a row of the periodic table.
  • The first ionization energy decreases as we go down a column of the periodic table.

The first trend isn't surprising. We might expect the first ionization energy to become larger as we go across a row of the periodic table because the force of attraction between the nucleus and an electron becomes larger as the number of protons in the nucleus of the atom becomes larger.

The second trend results from the fact that the principal quantum number of the orbital holding the outermost electron becomes larger as we go down a column of the periodic table. Although the number of protons in the nucleus also becomes larger, the electrons in smaller shells and subshells tend to screen the outermost electron from some of the force of attraction of the nucleus. Furthermore, the electron being removed when the first ionization energy is measured spends less of its time near the nucleus of the atom, and it therefore takes less energy to remove this electron from the atom.

The figure below shows the first ionization energies for elements in the second row of the periodic table. Although there is a general trend toward an increase in the first ionization energy as we go from left to right across this row, there are two minor inversions in this pattern. The first ionization energy of boron is smaller than beryllium, and the first ionization energy of oxygen is smaller than nitrogen.

These observations can be explained by looking at the electron configurations of these elements. The electron removed when a beryllium atom is ionized comes from the 2s orbital, but a 2p electron is removed when boron is ionized.

Be: [He] 2s2

B: [He] 2s2 2p1

The electrons removed when nitrogen and oxygen are ionized also come from 2p orbitals.

Ionization Energy Of Fluorine

N: [He] 2s2 2p3

O: [He] 2s2 2p4


But there is an important difference in the way electrons are distributed in these atoms. Hund's rules predict that the three electrons in the 2p orbitals of a nitrogen atom all have the same spin, but electrons are paired in one of the 2p orbitals on an oxygen atom.

Hund's rules can be understood by assuming that electrons try to stay as far apart as possible to minimize the force of repulsion between these particles. The three electrons in the 2p orbitals on nitrogen therefore enter different orbitals with their spins aligned in the same direction. In oxygen, two electrons must occupy one of the 2p orbitals. The force of repulsion between these electrons is minimized to some extent by pairing the electrons. There is still some residual repulsion between these electrons, however, which makes it slightly easier to remove an electron from a neutral oxygen atom than we would expect from the number of protons in the nucleus of the atom.

Practice Problem 4:

Predict which element in each of the following pairs has the larger first ionization energy.

(a) Na or Mg

(b) Mg or Al

(b) Mg or Al

(c) F or Cl

By now you know that sodium forms Na+ ions, magnesium forms Mg2+ ions, and aluminum forms Al3+ ions. But have you ever wondered why sodium doesn't form Na2+ ions, or even Na3+ ions? The answer can be obtained from data for the second, third, and higher ionization energies of the element.

The first ionization energy of sodium, for example, is the energy it takes to remove one electron from a neutral atom.

Na(g) + energy Na+(g) + e-

The second ionization energy is the energy it takes to remove another electron to form an Na2+ ion in the gas phase.

Na+(g) + energy Na2+(g) + e-

The third ionization energy can be represented by the following equation.

Na2+(g) + energy Na3+(g) + e-

The energy required to form a Na3+ ion in the gas phase is the sum of the first, second, and third ionization energies of the element.

First, Second, Third, and Fourth Ionization Energies
of Sodium, Magnesium, and Aluminum (kJ/mol)

It doesn't take much energy to remove one electron from a sodium atom to form an Na+ ion with a filled-shell electron configuration. Once this is done, however, it takes almost 10 times as much energy to break into this filled-shell configuration to remove a second electron. Because it takes more energy to remove the second electron than is given off in any chemical reaction, sodium can react with other elements to form compounds that contain Na+ ions but not Na2+ or Na3+ ions.

A similar pattern is observed when the ionization energies of magnesium are analyzed. The first ionization energy of magnesium is larger than sodium because magnesium has one more proton in its nucleus to hold on to the electrons in the 3s orbital.

Mg: [Ne] 3s2

The second ionization energy of Mg is larger than the first because it always takes more energy to remove an electron from a positively charged ion than from a neutral atom. The third ionization energy of magnesium is enormous, however, because the Mg2+ ion has a filled-shell electron configuration.

The same pattern can be seen in the ionization energies of aluminum. The first ionization energy of aluminum is smaller than magnesium. The second ionization energy of aluminum is larger than the first, and the third ionization energy is even larger. Although it takes a considerable amount of energy to remove three electrons from an aluminum atom to form an Al3+ ion, the energy needed to break into the filled-shell configuration of the Al3+ ion is astronomical. Thus, it would be a mistake to look for an Al4+ ion as the product of a chemical reaction.

Practice Problem 5:

Predict the group in the periodic table in which an element with the following ionization energies would most likely be found.

1st IE = 786 kJ/mol

2nd IE = 1577

3rd IE = 3232

4th IE = 4355

5th IE = 16,091

6th IE = 19,784

Practice Problem 6:

Use the trends in the ionization energies of the elements to explain the following observations.

(a) Elements on the left side of the periodic table are more likely than those on the right to form positive ions.

(b) The maximum positive charge on an ion is equal to the group number of the element

Ionization energies measure the tendency of a neutral atom to resist the loss of electrons. It takes a considerable amount of energy, for example, to remove an electron from a neutral fluorine atom to form a positively charged ion.

The electron affinity of an element is the energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion. A fluorine atom in the gas phase, for example, gives off energy when it gains an electron to form a fluoride ion.

F(g) + e- F-(g)Ho = -328.0 kJ/mol

Electron affinities are more difficult to measure than ionization energies and are usually known to fewer significant figures. The electron affinities of the main group elements are shown in the figure below.

Several patterns can be found in these data.

  • Electron affinities generally become smaller as we go down a column of the periodic table for two reasons. First, the electron being added to the atom is placed in larger orbitals, where it spends less time near the nucleus of the atom. Second, the number of electrons on an atom increases as we go down a column, so the force of repulsion between the electron being added and the electrons already present on a neutral atom becomes larger.
  • Electron affinity data are complicated by the fact that the repulsion between the electron being added to the atom and the electrons already present on the atom depends on the volume of the atom. Among the nonmetals in Groups VIA and VIIA, this force of repulsion is largest for the very smallest atoms in these columns: oxygen and fluorine. As a result, these elements have a smaller electron affinity than the elements below them in these columns as shown in the figure below. From that point on, however, the electron affinities decrease as we continue down these columns.

At first glance, there appears to be no pattern in electron affinity across a row of the periodic table, as shown in the figure below.

Ionization Energy Of Fluorine

When these data are listed along with the electron configurations of these elements, however, they make sense. These data can be explained by noting that electron affinities are much smaller than ionization energies. As a result, elements such as helium, beryllium, nitrogen, and neon, which have unusually stable electron configurations, have such small affinities for extra electrons that no energy is given off when a neutral atom of these elements picks up an electron. These configurations are so stable that it actually takes energy to force one of these elements to pick up an extra electron to form a negative ion.

Electron Affinities and Electron Configurations for the First 10 Elements in the Periodic Table

Element Electron Affinity (kJ/mol)Electron Configuration
H72.8 1s1
He<0 1s2
Li59.8[He] 2s1
Be <0[He] 2s2
B27[He] 2s2 2p1
C122.3[He] 2s2 2p2
N<0[He] 2s2 2p3
O 141.1 [He] 2s2 2p4
F328.0[He] 2s2 2p5
Ne<0[He] 2s2 2p6

Students often believe that sodium reacts with chlorine to form Na+ and Cl- ions because chlorine atoms 'like' electrons more than sodium atoms do. There is no doubt that sodium reacts vigorously with chlorine to form NaCl.

2 Na(s) + Cl2(g) 2 NaCl(s)

Furthermore, the ease with which solutions of NaCl in water conduct electricity is evidence for the fact that the product of this reaction is a salt, which contains Na+ and Cl- ions.

Ionization Energy Of Fluorine Oxygen

The only question is whether it is legitimate to assume that this reaction occurs because chlorine atoms 'like' electrons more than sodium atoms.

What Is The Ionization Energy Of Fluorine

The first ionization energy for sodium is one and one-half times larger than the electron affinity for chlorine.

Na: 1st IE = 495.8 kJ/mol

Cl: EA = 328.8 kJ/mol

Thus, it takes more energy to remove an electron from a neutral sodium atom than is given off when the electron is picked up by a neutral chlorine atom. We will obviously have to find another explanation for why sodium reacts with chlorine to form NaCl. Before we can do this, however, we need to know more about the chemistry of ionic compounds.

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