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Learning objectives
- Relate ionization energies to elemental chemistry
We have seen that when elements react, they often gain or lose enough electrons to reach the valence electron configuration of the nearest noble gas. In this section, we will develop a more quantitative approach to predicting such reactions by examining the periodic trends of energy changes that accompany ion formation.
Ionization activities
Since atoms do not spontaneously lose electrons, energy is required to remove an electron from an atom to form a cation. Chemists define the ionization energy (\(I\)) of an element as the amount of energy needed to remove an electron from a gaseous \(E\) atom in its ground state. \(I\) is therefore the energy needed for the reaction
\[ E_{(g)} \right arrow E^+_{(g)} +e^- \;\;\ \text{action required=I } \label{9.4.1}\]
Since an energy input is required, the ionization energy is always positive (\(I > 0\)) for the reaction, as written in Equation 9.4.1. Its higher valuesImeans the electron is more tightly bound to the atom and harder to remove. The standard units of ionization energy are kilojoules/mole (kJ/mol) or electronvolts (eV):
\[1\; eV/atom = 96.49\; kJ/mol]
If an atom has more than one electron, the amount of energy needed to remove more electrons increases steadily. We can define the first ionization energy (I1), the second ionization energy (I2) and generally oneNionization energy (IN) according to the following reactions:
\[ E_{(g)} \right arrow E^+_{(g)} +e^- \;\;\ I_1=\text{1. ionization energy} \label{9.4.2}\]
\[ E_{(g)} \right arrow E^+_{(g)} +e^- \;\;\ I_2=\text{2. ionization energy} \label{9.4.3}\]
\[ E^+_{(g)} \rightarrow E^{2+}_{(g)} +e^- \;\;\ I_3=\text{3. ionization energy} \label{9.4.4} \]
The values of the ionization energies \(Li\) and \(Be\) given in the table \(\PageIndex{1}\) show that successive ionization energies for an element are constantly increasing. This means that removing the second electron from an atom requires more energy than removing the first electron, and so on. There are two reasons for this trend. First, the second electron is removed from the positively charged particle, not from the neutral one, so according to Coulomb's law, more energy is needed. Second, the removal of the first electron reduces the repulsive forces between the remaining electrons, so the attraction of the remaining electrons to the nucleus is stronger.
Note
Successive ionization energies for an elementincreasestable.
| Reaction | Electronic transition | \[I\] | Reaction | Electronic transition | \[I\] |
|---|---|---|---|---|---|
| \(Li_{(g)}\rightarrow Li^+_{(g)} + e^-\) | \(1s^22s^1 \right arrow 1s^2\) | I1= 520,2 | \(Be_{(g)} \rightarrow Be^+_{(g)}+e^-\) | \(1s^22s^2 \right arrow 1s^22s^1\) | I1= 899,5 |
| \(Li^+_{(g)} \rightarrow Li^{2+}_{(g)} +e^-\) | \(1s^2 \right arrow 1s^1\) | I2= 7298,2 | \(Be^+_{(g)} \rightarrow Be^{2+}_{(g)} +e^-\) | \(1s^22s^1 \right arrow 1s^2\) | I2= 1757,1 |
| \(Li^{2+}_{(g)} \rightarrow Li^{3+}_{(g)} + e^-\) | \(1s^1 \right arrow 1s^0\) | I3= 11,815,0 | \(Be^{2+}_{(g)} \rightarrow Be^{3+}_{(g)}+e^-\) | \(1s^2 \right arrow 1s^1\) | I3= 14,848,8 |
| \(Be^{3+}_{(g)} \rightarrow Be^{4+}_{(g)}+e^-\) | \(1s^1 \right arrow 1s^0\) | I4= 21,006,6 |
The most important consequence of the values given in the table \(\PageIndex{1}\) is that the chemical composition \(Li\) is dominated by the ion \(Li^+\), while the chemical composition \(Be\) ) is dominated by by the +2 oxidation state. Action required to remove itotherelectron from \(Li\)
\[Li^+_{(g)} \rightarrow Li^{2+}_{(g)} + e^− \label{9.4.5}\]
is more than 10 times the energy needed to remove the first electron. So is the energy needed to remove itthirdelectron from \(Be\)
\[Be^{2+}_{(g)} \rightarrow Be^{3+}_{(g)} + e^− \label{9.4.6}\]
it is about 15 times the energy needed to remove the first electron and about 8 times the energy needed to remove the second electron. Both \(Li^+\) and \(Be^{2+}\) have onessmall2closed-shell configurations, and much more energy is required to remove an electron from 1small2core than with 2smallvalence orbital of the same element. The chemical implications are enormous: lithium (and all alkali metals) forms compounds with the 1+ ion, but not with the 2+ or 3+ ions. Similarly, beryllium (and all alkaline earth metals) forms compounds with the 2+ ion, but not with the 3+ or 4+ ions.The energy needed to remove electrons from a full nucleus is too much and simply cannot be achieved in normal chemical reactions.
Note
The energy required to remove electrons from a filled nucleus is too great under normal reaction conditions.
Ionizing effectsmall- IPI number-Lock items
The ionization energies of elements in the third row of the periodic table follow the same pattern as for \(Li\) and \(Be\) (Table \(\PageIndex{2}\)): successive ionization energies increase steadily as Electrons are removed from valence orbitals (3smallthe 3PI number, in this case) followed by a particularly large increase in ionization energy as electrons are removed from filled core levels, as indicated by the bold diagonal line in the table \(\PageIndex{2}\). Thus, in the third row of the periodic table, the greatest increase in ionization energy corresponds to the removal of the fourth electron from \(Al\), the fifth electron from Si, and so on, i.e. the removal of an electron from an ion that has the valence electron configuration of the previous noble gas. This formula explains why elemental chemistry usually involves only valence electrons. Too much energy is required to remove or share internal electrons.
| Element | I 1 | I 2 | I 3 | I 4 | I 5 | I 6 | I 7 |
|---|---|---|---|---|---|---|---|
| * Inner shell electron | |||||||
| Already | 495,8 | 4562,4* | — | — | — | — | — |
| Mg | 737,7 | 1450,7 | 7732,7 | — | — | — | — |
| Lap | 579.4.4 | 1816.7 | 2744,8 | 11.579.4.4 | — | — | — |
| I | 786,5 | 1577.1 | 3231,6 | 4355,5 | 16.090,6 | — | — |
| PI number | 1011.8 | 1909.4.4 | 2914.1 | 4963.6 | 6274,0 | 21.269.4.3 | — |
| small | 999,6 | 2251,8 | 3357 | 4556.2 | 7004.3 | 8495,8 | 27109.4.3 |
| Kl | 1251.2 | 2297,7 | 3822 | 5158,6 | 6540 | 9362 | 11.018,2 |
| Ar | 1520,6 | 2665,9 | 3931 | 5771 | 7238 | 8781,0 | 11.995,3 |
Example \(\PageIndex{1}\): Highest fourth ionization energy
Based on their position in the periodic table, predict which of these elements has the highest fourth ionization energy: B, C, or N.
Dany:three items
Ask:the element with the fourth highest ionization energy
Strategy:
- List the electron configuration of each element.
- Determine whether electrons are removed from a filled or partially filled valence shell. Guess which element has the highest fourth ionization energy, considering that the highest energy corresponds to the removal of electrons from an electron-filled nucleus.
Solution:
ONEAll of these elements are in the second row of the periodic table and have the following electron configurations:
- B: [On] 2small22PI number1
- DO:[On] 2small22PI number2
- N: [On] 2small22PI number3
andThe fourth ionization energy of an element (\(I_4\)) is defined as the energy needed to remove the fourth electron:
\[E^{3+}_{(g)} \rightarrow E^{4+}_{(g)} + e^-\]
Since carbon and nitrogen have four and five valence electrons respectively, their fourth ionization energy corresponds to the removal of an electron from a partially filled valence shell. In contrast, the fourth ionization energy for boron corresponds to the removal of an electron from the filled 1small2subshell. This should require much more energy. Actual values are: B, 25,026 kJ/mol; C, 6223 kJ/mol; and N, 7475 kJ/mol.
Exercise \(\PageIndex{1}\): Lowest second ionization energy
Based on their position in the periodic table, determine which of these elements has the lowest second ionization energy: Sr, Rb, or Ar.
Answer:Sr
The first column of data in the table \(\PageIndex{2}\) shows that the first ionization energies tend to increase along the third row of the periodic table. This is because the valence electrons do not shield very well, allowing the effective nuclear charge to increase steadily throughout the series. Therefore, the valence electrons are more strongly attracted to the nucleus, so the size of the atoms decreases and the ionization energies increase. These phenomena represent two sides of the same coin: stronger electrostatic interactions between the electrons and the nucleus further increase the energy needed to remove the electrons.
However, the first ionization energy decreases in Al ([Ne]3small23PI number1) iw N ([Ne]3small23PI number4). Electrons filled with aluminum 3small2The subshell is better in test 3PI number1electron than is controlled by the nuclear charge, sosmallelectrons penetrate closer to the nucleus thanPI numberdoes the electron. The reduction in S occurs because the two electrons are in the samePI numberorbitally repel each other. This makes the S atom a bit less stable than you might expect, as is the case with all group 16 elements.
Image \(\PageIndex{2}\): First ionization activitiessmall-,PI number-,Hello-, Ieat-Lock items
The first ionization energies of elements in the first six rows of the periodic table are shown in the figure \(\PageIndex{1}\) and numerically and graphically in the figure \(\PageIndex{2}\). These figures illustrate three important trends:
- Changes observed in the second (Li to Ne), fourth (K to Kr), fifth (Rb to Xe) and sixth (Cs to Rn) seriessmallIPI numberthe blocks have a pattern similar to that described for the third row of the periodic table. However, the transition metals are contained in the fourth, fifth and sixth rows, and the lanthanides in the sixth row. The first ionization energies of the transition metals are somewhat similar, as are the energies of the lanthanides. Ionization energies increase from left to right in each row, with deviations shown inns2np1(group 13),ns2np4(group 16) ins2(N− 1)Hello10(group 12) electron configurations.
- The first ionization energies generally decrease in the column. Although the main quantum noNincreases in the column, the filled inner shells are effective in shielding the valence electrons, so there is a relatively small increase in the effective nuclear charge. Consequently, atoms become larger as they gain electrons. Valence electrons farther from the nucleus are less tightly bound, making them easier to remove, resulting in a decrease in ionization energy.A larger radius corresponds to a lower ionization energy.
- Due to the first two trends, the elements that form positive ions more easily (have the lowest ionization energies) are in the lower left corner of the periodic table, while those that are more difficult to ionize are in the upper right corner of the periodic table. Consequently, ionization energies generally increase diagonally from lower left (Cs) to upper right (He).
Note
In general, \(I_1\) increases diagonally from the lower left corner of the periodic table to the upper right corner.
The darkness of the shading in the table cells indicates the relative magnitudes of the ionization energies. Elements marked in gray have undetermined first ionization energies. Source: Data fromCRC Handbook of Chemistry and Physics(2004).
Gallium (Ga), which is the first element after the first series of transition metals, has the following electron configuration: [Ar]4small23Hello104PI number1. Its first ionization energy is much lower than that of the immediately preceding element, zinc, because the filling of the 3rdHello10The gallium subshell is inside 4PI numbersubshell, sort of single 4PI numberelectron from the nucleus. The experiments revealed something even more interesting: the second and third electrons removed during gallium ionization come from the 4th electron.small2path,NO3Hello10subshell. Gallium chemistry is dominated by the formed Ga3+ion with its [Ar]3Hello10electron configuration. This and similar electron configurations are particularly stable and are often found in heavier onesPI number-block items. They are sometimes referred to as pseudo-noble gas configurations. In fact, for elements exhibiting these configurations,There are no known chemical compounds where electrons are removed from (n - 1)Hello10 full subshell.
As we have seen, the first ionization energies of the transition metals and lanthanides vary very little in each series. The differences in their second and third ionization energies are also rather slight, in sharp contrast to the observed patternsmall- IPI number-block items. The reason for these similarities is that transition metals and lanthanides form cations through lossnselectrons before (N− 1)Hello(N− 2)eatelectrons respectively. This means that the transition metal cations have (N− 1)HelloNthe valence electron configurations and the lanthanide cations have (N− 2)eatNvalence electron configurations. Because (N− 1)HelloI (N− 2)eatthe crusts are closer to the core thannsshell, (N− 1)HelloI (N− 2)eatelectrons scan itnselectrons quite effectively, reducing the effective nuclear charge felt by onselectrons. HowGincreases, the increasing positive charge is largely canceled out by the electrons added to (N− 1)Hello(N− 2)eatorbital.
Thatnselectrons are removed before (N− 1)Hello(N− 2)eatThe electrons may surprise you because the orbitals were filled in reverse order. Actually,ns, (N− 1)Hello, I (N− 2)eatThe orbitals are so close together in energy and interpenetrate so widely that very small changes in the effective nuclear charge can change the order of their energy levels. asHelloorbitals fill up, active nuclear charge causes 3Hellothe orbitals are slightly lower in energy than 4smallorbitals. [Would] 3Hello2electronic configuration of Ti2+tells us that 4smallTitanium electrons are lost before 3Helloelectrons; this is confirmed by the experiment. A similar pattern is observed with the lanthanides, which produce cations from (N− 2)eatNvalence electron configuration.
Because their first, second, and third ionization energies change so little over a series, these numbers are importanthorizontalsimilarities in chemical properties other than the expected vertical similarities. For example, all first-order transition metals except scandium form stable compounds such as M2+ions, while the lanthanides mainly form compounds where they occur as M3+ions.
Example \(\PageIndex{2}\): Decrease the first ionization energy
Use their position in the periodic table to predict which element has the lowest first ionization energy: Ca, K, Mg, Na, Rb or Sr.
Dany:six positions
Ask:the element with the lowest first ionization energy
Strategy:
Locate the elements in the periodic table. From the ionization energy trends up the row and down the column, identify the element with the lowest first ionization energy.
Solution:
These six elements form a rectangle on the leftmost two columns of the periodic table. Since we know that the ionization energies increase from left to right in a row and from bottom to top of a column, we can predict that the element in the lower left corner of the rectangle will have the lowest first ionization energy: Rb.
Exercise \(\PageIndex{2}\): Higher first ionization energy
Use their position in the periodic table to predict which element has the highest first ionization energy: As, Bi, Ge, Pb, Sb or Sn.
Answer:How
Ionization energy:https://youtu.be/k7j-u02ifzo
summary
- In general, the first values of ionization energies and electronegativity increase diagonally from the lower left of the periodic table to the upper right, and electron affinities become more negative in series.
The tendency of an element to lose is one of the most important factors determining the type of compounds it creates. Periodic behavior is most evident forionization energy (I), the energy needed to remove an electron from a gaseous atom. The energy required to remove more electrons from an atom increases steadily, with a significant increase occurring as an electron is removed from a filled inner shell. As a result, only valence electrons can be removed by chemical reactions, leaving the filled inner shell intact. Ionization energies explain the common oxidation states observed for the elements. Ionization energies increase diagonally from the lower left corner of the periodic table to the upper right corner. Small deviations from this trend can be explained by particularly stable electronic configurations, the so-calledconfigurations of pseudo-noble gases, either in the parent atom or in the resulting ion.
FAQs
What is ionization energy 1 and 2? ›
The symbol I1 stands for the first ionization energy (energy required to take away an electron from a neutral atom) and the symbol I2 stands for the second ionization energy (energy required to take away an electron from an atom with a +1 charge. Each succeeding ionization energy is larger than the preceding energy.
What is 1st 2nd and 3rd ionization enthalpy? ›First, second and third ionization energies are 737, 1045 and 7733 kJ/mol respectively.
What is meant by 1 ionization energy? ›The first ionization energy is the energy required to remove the most loosely held electron from one mole of neutral gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+. This is more easily seen in symbol terms. X(g)→X+(g)+e− It is the energy needed to carry out this change per mole of X.
What is 2 ionization energy? ›An element's second ionization energy is the energy required to remove the outermost, or least bound, electron from a 1+ ion of the element. Because positive charge binds electrons more strongly, the second ionization energy of an element is always higher than the first. Created by Jay.
What are the 2 types of ionization energy? ›Ionization of molecules often leads to changes in molecular geometry, and two types of (first) ionization energy are defined – adiabatic and vertical.
What is the ionization energy of each element? ›| Ionization Energy | Name chemical element | Symbol |
|---|---|---|
| 7,8994 | Germanium | Ge |
| 7,9024 | Iron | Fe |
| 8,1517 | Silicon | Si |
| 8,298 | Boron | B |
The second ionization energy is always larger than the first ionization energy, because it requires even more energy to remove an electron from a cation than it is from a neutral atom.
What is first and second ionization energy examples? ›The first ionization energy of sodium, for example, is the energy it takes to remove one electron from a neutral atom. The second ionization energy is the energy it takes to remove another electron to form an Na2+ ion in the gas phase.
What is the highest first ionization energy? ›Nitrogen has the highest first ionization energy, though it should have less than that of oxygen while going across a period. The electronic configuration of nitrogen is 1s2,2s2,2px1,2py1,2pz1. Nitrogen has three electrons in 2p subshell which are singly filled.
What is the first ionisation energy equation? ›1st ionisation energy: K(g) → K+(g) + e– n=1 2nd ionisation energy: K+(g) → K2+(g) + e– n=2 Your task 1.
What is the order of the ionization energy? ›
Explanation: Since it is easier to take out an electron from a negatively charged ion than from a neutral atom, the ionization energies for anions will be lesser than the corresponding neutral atoms. So the order of ionisation energy is F>F− and Cl>Cl−.
Why is ionization energy? ›An element's first ionization energy is the energy required to remove the outermost, or least bound, electron from a neutral atom of the element. On the periodic table, first ionization energy generally increases as you move left to right across a period.
What is 2nd 3rd ionization energy? ›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.
What is 1st ionisation enthalpy 2nd ionization enthalpy? ›Hint: First ionization enthalpy is the energy required to remove the first most loosely bound electron from an isolated atom in its ground state while second ionization enthalpy is the energy required to remove the second most loosely bound electron.
Which element has the lowest 1st ionization energy? ›Thus lowest IE in this case is of Bi as it is bottom most element.
What does a low ionization energy mean? ›If the ionization energy is low, that means that it takes only a small amount of energy to remove the outermost electron.
What are examples of ionization? ›In chemistry, ionization often occurs in a liquid solution. For example, neutral molecules of hydrogen chloride gas, HCl, react with similarly polar water molecules, H2O, to produce positive hydronium ions, H3O+, and negative chloride ions, Cl-.
Which period 2 has the highest ionization energy? ›Neon (Ne) is present at the rightmost side in Period 2 and therefore, it has the maximum ionization energy in the second Period.
How do you explain ionization? ›Ionization is the process by which ions are formed by gain or loss of an electron from an atom or molecule. If an atom or molecule gains an electron, it becomes negatively charged (an anion), and if it loses an electron, it becomes positively charged (a cation). Energy may be lost or gained in the formation of an ion.
Which group has the highest ionization energy? ›Therefore, noble gases will have the highest ionisation energy.
Which of the following has lowest ionization energy? ›
Therefore, the correct option is (C) Sulfur.
Why does ionization energy decrease down a group? ›On the periodic table, first ionization energy generally decreases as you move down a group. This is because the outermost electron is, on average, farther from the nucleus, meaning it is held less tightly and requires less energy to remove.
What is the difference between the first and second ionization energy? ›The first ionization energy is the energy required to remove the 1st electron, the second ionization energy is the energy to remove the second electron, and so on and so forth. Moving left to right within a period or upward within a group, the first ionization energy generally increases.
What is higher ionization energy? ›The higher the ionization energy, the more difficult it is to remove an electron. Therefore, ionization energy is in indicator of reactivity. Ionization energy is important because it can be used to help predict the strength of chemical bonds.
Which has the highest difference in first ionization energy? ›Thus, the element having the greatest difference between first and second ionization energies is K.
Which element has the lowest second ionization energy? ›Thus, the second ionization energy of Beryllium is the lowest ionization energy.
What are the first five ionization energies? ›The first five ionization energies of an element are 9.1, 16.2, 24.5, 35 and 205.7 eV respectively. The number of valence electrons in the atom is. No worries! We've got your back.
Which element has the highest first ionization energy Why? ›Fluorine has highest first ionization energy since it has the tendency to gain electron and no tendency to lose the electron from its valence shell.
Why is 3rd ionization energy higher than 1st? ›On removal of 2 electrons from Ca, it attains configuration of Ar which is an inert gas and its configuration is extremely stable. Hence excessive energy is needed to break this configuration. Thus the third IE of Ca is significantly higher than that of its second or first.
What affects ionization energy? ›Ionization energy depends on two factors: The force of attraction between electrons and the nucleus. The force of repulsion between electrons.
Does ionization increase down a group? ›
Ionization energy (IE) is the energy required to remove the highest-energy electron from a neutral atom. In general, ionization energy increases across a period and decreases down a group. Across a period, effective nuclear charge increases as electron shielding remains constant.
What is the difference between the first and second ionization energy of Group 1 and 2 elements? ›Definition: Second Ionization Energy
Group 1 elements have first ionization energies that are much smaller than their second ionization energies, and group 2 elements have second ionization energies that are much smaller than their third ionization energies.
The first ionization energy of the oxygen molecule is the energy required for the following process: O21g2¡O2 +1g2 + e- The energy needed for this process is 1175 kJ>mol, very similar to the first ionization energy of Xe.
Why is the 3rd ionization energy greater than the 2nd? ›The third ionization energy is even higher than the second. Successive ionization energies increase in magnitude because the number of electrons, which cause repulsion, steadily decrease. This is not a smooth curve There is a big jump in ionization energy after the atom has lost its valence electrons.
How do you calculate second ionization energy? ›The equation En=-Z2RH/n2 gives the energy (according to the Bohr model) of the energy levels available to an electron in any hydrogen-like atom.
Which element has the highest 2nd ionization energy? ›Hence, second ionization energy of sodium is largest.
Does second ionisation energy increase down a group? ›Going down a group, the ionisation energy decreases.
What is the difference between first ionization energy and second ionization energy quizlet? ›The first ionization energy is the energy needed to remove a first electron from an atom. The second ionization energy is the energy to remove a second electron.
Why O2 has lowest ionization energy? ›O2− has lower ionization energy than any other. Since it has contained more numbers of electrons in its molecular orbital hence the effective nuclear charge was decreased. So, it has low ionization energy.
Does O or O2 have higher ionization energy? ›Atomic oxygen has higher ionization energy than oxygen molecules due to the repulsive force of electrons in molecular oxygen.
What is the trend of ionization energy? ›
What is the trend for ionization energy? The general trend of ionization energy in a periodic table is that the energy increases as you move in the direction of hydrogen to helium and decreases as you move from hydrogen to cesium.
What is the difference between first and second ionization energy? ›First ionization energy: The energy needed to remove the outermost, or highest energy, electron from a neutral atom in the gas phase. Second ionization energy: The energy it takes to remove an electron from a 1+ ion (meaning the atom has already lost one electron and now removing the second).
What does ionization energy depend on? ›Ionization energy depends on orbital energy, which depends on the type of orbital and the effective nuclear charge. Thus, it follows predictable patterns in the periodic table. As you go down, n increases, and the energy of the orbital increases.
Which element has the lowest first ionization energy? ›The value of ionization energy (IE) decreases down the group due to the increasing size as the valence electrons are more loosely bound. Thus lowest IE in this case is of Bi as it is bottom most element.