Periodic Property: Electron Gain Enthalpy or Electron Affinity

You must have heard about energy. Either something is being done or happening spontaneously, energy is involved in every process. Whenever you talk about a process you talk about the energy changes. Just like when you are talking about a business, you talk about profit and loss, the energy change in every process is also in the terms of profit and loss. Similar to the financial profit and loss, energy is also influenced by various factors and we have to keep these factors constant to get the accurate amount of energy. Pressure and volume are two variables for this process so if you keep them constant, you can get the accurate value of energy for this process. This energy is termed as the enthalpy of the process or in the other words, if a process is carried out at constant pressure and volume, the energy change is called the enthalpy of the process. 

Electron Gain Enthalpy is opposite to the Ionization enthalpy. As you know that Ionization Enthalpy is the energy required to remove an electron from a neutral atom. On the contrary, when an electron is added to a neutral atom, a certain amount of energy is required or released. This energy is termed as Electron Gain Enthalpy.  

When an electron is added to an atom it gets support from the nucleus in the form of nuclear attraction but it also has to face repulsion from other electrons. That is why Electron Gain Enthalpy depends on the element’s nature as well as its requirements. If an element is going to complete its octet on adding an electron, it will accept the incoming electron easily. When it is beneficial for an element to add an electron, it requires lesser energy to do so or even sometimes energy is released. When addition of electron requires energy, EGE  gets a positive sign and when it releases energy, EGE gets a negative sign.

Whenever you talk about any property/quality of an element you have to consider all the specifications of its atom such as size of the atom, nuclear charge, amount of shielding, electronic arrangement and the type of electron which is involved in the process. Let’s see how these factors affect the electron gain enthalpy.

• Size of the atom: In a smaller atom electrons are placed closely to the nucleus, so the nucleus can attract the incoming electron and requires lesser energy to accommodate an extra electron. 
• Nuclear charge: Larger nuclear charge supports the process by attracting the incoming electron more efficiently.
• Shielding effect of inner sub-shells: When the electron is added to the outer sub-shell, it will be favorable if the inner sub-shells have weak shielding effect so that the incoming electron can seek out support from the nucleus and experience better nuclear attraction. (order of shielding effect s > p > d > f).
• The type of sub-shell where the incoming electron is being added (s, p, d or f): Electrons are packed most tightly in s sub-shell as it is the smallest one. Sub-shell p has more space for electrons than s, so it will be easier to insert an extra electron in p sub-shell than in s sub-shell. In a larger sub-shell incoming electron has to face lesser repulsion from other electrons. The order of the electron gain enthalpy for the sub-shells is s > p > d > f.
• Electronic configuration: As we discussed earlier that half-filled and full-filled electronic configurations have extra stability. So the atoms that need an electron to achieve half-filled or full-filled state of its outer most sub-shell easily accommodate incoming electron and require lesser energy.

As we go downward in a column the size of atom increases, so it would be difficult for nucleus to attract incoming electron at a distance. So the addition is to be done forcefully, which means certain amount of energy is required. That’s why elements towards the bottom of column have low negative values of Electron Gain Enthalpy (release lesser amount of energy)  as compared to the elements towards the top.

Elements of 1^stand 2^nd group (column) have least negative electron gain enthalpy. They can achieve octet on losing electron so they prefer to lose electron rather than gaining it. Lower negative values show that they accept the incoming electron quite unhappily.

Noble elements of 18^th  group have quite large positive values of electron gain enthalpy. It means that the electron is being introduced forcefully and they have to accommodate it unwillingly, as a result of which a large amount of energy is required.

As we go towards right side in a period (row), the size of the atom decreases and the strength of the nucleus (nuclear charge) increases. Since the addition of electron is supported by the strength of nucleus, the process is accomplished easily with lesser energy.

General Trend of Electron Gain Enthalpy

Elements of 17^th  group have larger negative values of electron gain enthalpy. If you write their electronic configuration you will find that they are just one electron short to achieve octet, that’s why their behaviour is very much welcoming for the incoming electron. They release large amount of energy on introducing an extra electron to them, which means the process is quite feasible.

The trend of electron gain enthalpy is not as symmetrical as ionization energy (enthalpy) because the factors deciding it are closely intedependent and and their effects on the process vary considerably.

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Periodic Property: Deviations in The Trend of Ionization energy

We discussed about ionization energy in the last post. In this post, we will discuss the factors affecting IE and will try to find out the reasons behind the unexpected behaviour of a few elements. There are a few deviations in the usual trend of IE, like

ü  ⁴Be has higher IE than ⁵ B. 

ü  ¹² Mg has higher IE than ¹³ Al

ü  Group 13 shows irregular trend of IE

ü  Elements of group 15 have higher IE than expected.

Let’s find out the reasons behind their remarkable behaviour. I hope you have understood the process and concept of ionization. So think about the factors that decide the amount of ionization energy. These factors are: 

• Size of the atom: In a smaller atom electrons are packed tightly as compared to a larger one. 
• Nucleus charge: Large nuclear charge holds electrons tightly.
• How effectively the inner sub-shells shield the outer electrons (shielding effect s > p > d > f). Strong shielding defends outer electrons from the nuclear attraction and makes the exit of outer electron easier, while weak shielding enables nucleus to attract outer electrons more powefully and to hold them tightly, consequently hindering the exit of electron.
• The type of sub-shell or electron is involved (s, p, d or f): s sub-shell placed closer to the nucleus than p and hold electrons tightly than p, similarly p is closer than d, and d is closer than f . That’s why the order of IE : s > p > d > f. Energy required for the removal of an electron belonging to s sub-shell is the highest and for the removal of an electron belonging to f sub-shell is the lowest.
• Electronic configuration: Half filled and fulfilled sub-shells have extra stability. That’s why extra energy is needed to break such configurations.

All these factors are interrelated. Let’s write down the electronic configuration of ⁴Be and ⁵B.

⁴Be : 1s² , 2s²

⁵B : 1s² , 2s² , 2p¹

In ⁴Be outer most electron belongs to the s sub-shell which is placed closer to the nucleus and holds the electron tightly. And s sub-shell is in its completely filled state, which is the most stable state, so it makes the exit of the outer most electron even more difficult. On the other hand, in ⁵B, the outer most electron belongs to the p sub-shell which is placed farther than the s sub-shell, and the electron is placed singly in an orbital, which is comparatively easier to remove.

And for the similar reason ¹² Mg has higher IE than ¹³ Al. You may check it yourself.

Let’s check the group 13 now. What is different here? It is placed just after the d block. Elements of this group show unexpected behaviour once in case of ³¹ Ga which has higher IE than ¹³ Al and secondly in case of ⁸¹ Tl which has higher IE than ⁴⁹ In.

If you notice the place of ³¹ Ga in periodic table and write the electronic configuration, you will find that it has a completely filled d sub-shell which shields its outer most single electron of psub-shell quite weakly, as a result of which the nucleus binds this electron more tightly and more energy is required to remove it.

Similarly ⁸¹ Tl also has a completely filled d and f sub-shells. Its single outer most electron of p sub-shell is even more weakly shielded by these d and f sub-shells, hence bound tightly by the nucleus. This makes its exit more difficult.

Now see the elements of group 15. Write the configuration and focus on the outer most electron. You will find that it is the third electron of p sub-shell. All orbitals of p sub-shell are singly occupied. This state of a sub-shell is called the half-filled state. It is the next stable state to the fully filled one. That’s why group 15 elements require more energy to remove their outer most electron from the half-filled p sub-shell.

The trend of second and third ionization energies are quite irregular. The reason behind this is the change in electronic configuration and effective nuclear charge resulting from  the removal of the first electron from the atom. This in turn changes major factors affecting the ionization energies.

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Periodic Property: Ionization energy

Now you are quite familiar with the atom and you know very well that the atom is neutral because it has equal number of electrons and protons. If we remove one of the electrons or add an extra electron to it, this disturbs the balance of charge and the atom no longer remains neutral. It becomes a charged species and is now called as “ion”. It either develops an excess of positive charge or negative charge.

You also know that electrons are arranged in orbits and sub-shells and every electron has an address. What makes an electron different from other? Energy! Yes, it creates all the difference. Energy of an electron of 1st orbit is lower than the electron of 2^nd orbit. If certain amount of energy is supplied to 1st orbit electron it can jump to 2^nd orbit. Similarly, on supplying large amount of energy (sufficient energy to overcome the nuclear attraction) to the outer most electron, it can be removed from the atom. This process is called ionization and the energy required to remove an outer most electron is called the ionization energy.

It is possible to remove more than one electrons from an atom but you have to proceed step by step.

            Energy required to remove an electron from neutral atom is called “First Ionization Energy”.

            Energy required to remove an electron from A⁺ is called “Second Ionization Energy”.

            Energy required to remove an electron from A⁺²  is called “Third Ionization Energy”.

Removal of successive electrons from an ion becomes more difficult because of increased nuclear charge. That’s why 2^nd IE is always larger than the 1^st IE, because in A⁺ ion electrons are bound more tightly due to increased nuclear charge. And for the similar reason 3^rd IE is larger than the 2^nd IE.

When you go downwards in a group (column), ionization energy decreases. Though the nuclear charge increases when you go downwards in a group, but it is easier to remove an electron from a larger atom (Remember! number of orbits increases, it means the size of atom or the atomic radius increases.)

When you go towards right in a row (period), ionization energy increases, because atomic size decreses in that direction. And it would be difficult to remove an electron from a smaller atom than from a larger atom, because electrons in a smaller atom are bound tightly as compared to the larger one.

In general IE decreases down the group and increases in a period. You will find that group 1 elements have lowest IE in their respective period. Noble gases have highest IE in their respective period.  There are few deviations in the trend though, like

ü  ⁴Be has higher IE than ⁵ B. 

ü  ¹² Mg has higher IE than ¹³ Al

ü  Group 13 shows irregular trend of IE

ü  Elements of group 15 have higher IE than expected.

In the next post we will find out the reasons behind their remarkable behaviour. 

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Periodic Property: Size of the atom/ atomic radius

Periodic Table is the address book of elements. Elements are arranged here in the increasing order of their atomic number. It keeps all the information about elements like size, nature, behaviour towards other element, strength and weakness. However these informations are hidden inside the periodic table and you have to learn, how to explore them. I will explain these properties one by one.

Size of The Atom: Across The Group

The secret of every quality lies in the subatomic particles of the atom and how these particles are arranged in an atom. Let’s examine the elements of first column (group). As you go down in a column, you will find the number of orbits increases, it means the size of atom or the atomic radius increases. The element that is placed lower in a column has larger atomic radius/ size than the element which is placed upper in a column.

Size of The Atom: Across The Row (Period)

As you move towards right in a row, you will find that the orbit number remains the same but the number of electrons and the number of protons increase. Nucleus becomes more powerful with the increasing number of protons and therefore it becomes capable of binding electrons more closely . That is why on going towards right side in a row, size of atom or the atomic radius decreases. An element placed towards right side of a row has smaller atomic size/ radius than the element placed on left side of the row.

Since atomic size decreases across the row/period, the elements of 1^st column are largest and the elements of 17^th column are the smallest. Elements of 18^th column are different because they already have a complete octets in their outermost orbit. So, they called noble elements and have noble characteristics. We will discuss them separately.

When you observe the trend of atomic radius across the row, you will find that it is not as steady as we expect. It decreases across the row but when you move from “d block” to “p block” there is a large contraction in size. And again after “f block” you will see larger contraction in the size of “ d block” elements.

Let’s find out the probable reasons for this unsteady trend in size? So, first note down the factors that are responsible for contraction in size.

• Nuclear charge/number of protons, which pulls electrons towards nucleus like a magnet.
• Distance between the outer electrons and nucleus, which prevents them to be pulled towards the nucleus.

Across the row, nuclear charge increases and ideally the distance between outer orbit and nucleus should decrease gradually. Nothing can change the nucleus charge and the distance between nucleus and orbit. But if something happens which increases or decreases the force experienced by electrons, it can change the trend.

Shielding effect 

 

The electrons of inner orbit cancel some of the nuclear charge, so that the electrons of outer orbit experience less attraction by nucleus. This is called shielding. The extent of shielding caused by different sub-shells is different because, every sub-shell has different shape.

Every sub-shell has few places where no electron is present, which is called node. These nodes act as  holes in a wall through which nucleus pulls electrons more strongly. In other words, presence of nodes weakens the shielding effect of sub-shell. The number of nodes can be guessed by the azimuthal quantum number (l).

s sub-shell is spherical and has l = 0,  means has no node so, it shields most strongly.

p sub-shell is dumbbell shaped and has  l = 1,  means has one node

d sub-shell is double dumbbell and has l = 2,  means two nodes

f sub-shell is complex shaped and has l = 3,  means three nodes, it shields most weakly.

The order of shielding effect of different sub-shells is as follows:
              s > p > d > f

That is why in 4th row the contraction in size across the row is larger, because extra electrons are added in the inner (n-1) d sub-shell. And d sub-shell doesn’t shield outer electrons effectively, so nucleus pulls the electrons more strongly and causes larger contraction in atomic size.

Similarly, in 6^th row when electrons are added to the (n-2) f sub-shell, which provides weakest shielding, it causes larger contraction in the size of d block elements after filling 14 electrons in f orbital (the f block elements). This is also named as lanthanide contraction after the name of first element (Lanthanum ⁵⁷La) of f block. Because of it the size of the elements of last two rows of “d block” becomes almost same even though a shell is added.

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Role of Quantum Numbers in Periodic Table

Periodic table is the most customised address book of elements. Elements are placed here in increasing order of their atomic number. Every element has got a specific place. As you know that every element has unique characteristics. So how is it possible to arrange them in rows and columns? You must have noticed the lengths of columns are different, and there are also a few incomplete rows and two rows are separately placed at the bottom of the table. Why there are 18 columns and 7 rows? In this post we will try to find some answers to these questions.

Rows in periodic table

Write the electronic configuration of first element of each row. Now find the principal quantum number for the outer most electron of each element. You will find “n” of the element is similar to the number of the row it belongs to.  

“s Block” 

When you look at the Periodic table, you will see there are two columns on the left side. Let’s write the electronic configuration of first two elements of the first column.

¹H = 1s¹

³Li = 1s², 2s¹

Notice the outer orbital of these elements; its “s” in both elements.

Now write the electronic configuration of first two elements of the second column.

⁴Be = 1s², 2s²

¹²Mg = 1s², 2s², 2p⁶, 3s²

Notice the outer orbital of these elements; its again “s” in both elements.

As you have noticed that, all the elements of these two columns have their last electron filled in “s” orbital. That’s why they are called “s Block” elements. Since “s” orbital can accommodate maximum 2 electrons that’s why “s block” comprises of two columns.

“p Block”

When you look on the right side of the table you will find 6 columns there. Now write the electronic configuration of at least first elements of each column.

In all these elements outer most electron is filled in the “p” orbital. That’s why these 6 columns are placed under “p Block”. You can guess why there are 6 columns in “p block”. Yes of course because “p” orbital can accommodate maximum 6 electrons. If you write the configuration of other elements of the same column you will find all of them have similar number of electrons in the “p” orbital. In other words, elements of the same column have same “l”, “m_l” and “m_s” quantum numbers.

“d Block”

You must have noticed the bridge of 10 columns in the periodic table. I know you guessed the right reason of the name of this block and reason behind the number of columns. Orbital d can accommodate 10 electrons that’s why it has 10 columns and each element has filled their outer most electron in “d” orbital.

Before you go to write the configuration, I want to tell you the (n+l) Rule. In the previous post about electronic configuration I have told you about energy order of the sub-shells.

Energy order of the sub-shells is as follows: s < p < d < f

When you compare the energy of sub-shells belonging to different orbits you will find

1s < 2s < 2p < 3s < 3p < 3d

The (n+l) rule, which I am going to explain, is about the energy of sub-shells. The sub-shell with higher (n+l) value has higher energy. Let’s check this rule in the energy order given above.

1s has n = 1, l = 0 so, (n+l) = 1

2s has n = 2, l = 0 so, (n+l) = 2

2p has n = 2, l = 1 so, (n+l) = 3

Yes this rule is working fine. Now workout yourself, the order of higher sub-shells 3d, 4s, 4p, 4d, and 4f.

3d has n = 3, l = 2 so, (n+l) = 5

4s has n = 4, l = 0 so, (n+l) = 4

So, the energy of 4s < energy of 3d.

4p has n = 4, l = 1 so, (n+l) = 5

There is tie between 3d and 4p. When there is a tie, the “Aufbau Principle” will decide which one has higher energy. According to this principle the orbit nearer to the nucleus has lower energy or in other words the orbit with higher “n” value has higher energy.

So, the energy of 4p > energy of 3d.

Now the complete energy order will be

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p

Like p block elements, elements of d block also has similar set of “l”, “m_l” and “m_s” quantum numbers.

Count the number of rows after which “d Block” starts. You will find that it starts from 4^th row. You can justify it from the energy order of sub-shells. Yes, because “d” sub-shell is not present in 1^st and 2^nd  orbit and electron cannot be filled in “3d” sub-shell before filling the “4s” sub-shell. That’s why row 1, 2 and 3 have a gap between “s block” and “p block”. And “d block” starts in 4^th row after “4s”.

“f Block”

It is shown as the extended part of the periodic table. In fact it is present between 6s and 5d but to maintain the symmetry of the periodic table it is placed separately. It has 14 columns because “f” sub-shell can accommodate maximum 14 electrons and the outer most electron of each element is filled in “f” orbital.

Do you remember, that 4^th orbit has 4sub-shells s, p, d and f? However in the energy order of sub-shells discussed above, we didn’t mention 4f sub-shell. Let’s find out its place in the energy order.

4f has n = 4, l = 3 so, (n+l) = 7

5p has n = 5, l = 1 so, (n+l) = 6

6s has n = 6, l = 0 so, (n+l) = 6

5d has n = 5, l = 2 so, (n+l) = 7

There is tie between 4f and 5d and 5p and 6s. In this situation, “Aufbau Principle” will be the tie breaker and 5d wins over 4f and 6s wins over 5p because greater “n” has greater energy.

Now the complete energy order will be

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p.

Position of ²He

After understanding the reasons behind the arrangement of blocks and the elements belonging to them, you will be able to guess the position of any element from its electronic configuration. When you write the configuration of ²He : 1s², you may place it in the “s Block” but, it is placed in “p Block”.

The arrangement of elements in perodic table also sets particular trends of their properties along the rows and columns such as size of atom, electonegativity, ionization enthalpy and electron gain enthalpy. These are known as periodic properties. In the coming posts I will discuss these properties and there trends one by one. 

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Quantum Numbers: Postal Address of an Electron

As you know, in an atom every electron has a particular place. Just as you have your postal address, every electron also has a particular address. Like a pin code, electrons have a set of Quantum Numbers. These quantum numbers give valuable information about an element. On the basis of quantum numbers of the outer most electron an element gets its particular place in the periodic table. In this post, I am going to explain the quantum numbers and in my next post I will correlate them to the periodic table.

Electronic Configuration of ⁵B

Write the configuration of ⁵B: 1s², 2s², 2p¹. Now arrange its electrons in the orbitals and spot the outer most electron. List down all the specifications you need to write the address of this electron.

1.      The orbit number
2.      The Sub-shell
3.   The orbital
4.      The orientation

You are correct; these 4 specifications are listed as quantum numbers and make the complete address of every element. I will explain these specifications one by one.

Orbit Number: Principal Quantum Number (n)

The principal quantum number tells about the orbit number. It is shown by “n”. Its value can be 1, 2, 3, 4, 5 …..

In the above example for the outer most electron, the value of “n” will be 2, as it is present in orbit number 2.

The Sub-shell: Azimuthalquantum number (l)

Azimuthal quantum number tells about the sub-shell. The symbol for it is “l”. Its value can be 0, 1, 2, 3, 4…(n-1). For any electron the value of “l” is always less than the value of “n”.

The value of “l” is assigned for every sub-shell. 

For “s” the assigned value of “l” is = 0

For sub-shell “p” the assigned value of “l” is = 1

For sub-shell “d” the assigned value of “l” is = 2

For sub-shell “f” the assigned value of “l” is = 3

In the above example for the outer most electron, the value of “l” will be 1, as it is placed in “p” sub-shell.

The Orbital: Magnetic quantum number (m_l)

Magnetic quantum number (m_l) tells about the particular orbital where electron is placed. As you know every sub-shell has different number of orbitals. These sub-shells have used the advance level of architecture, so that every orbital has a definite orientation in the space. As in “p” sub-shell, one orbital is oriented in “x” axis, other one is in “y” axis and the third one is in the “z” axis. They named as (p_x ) _, (p_y) and (p_z )   according to their orientation in the space.

Magnetic quantum number is the notation for that particular orbital. Like in car we refer different seats as driver seat, window seat, front seat, middle seat and back seat.

The value of Magnetic quantum number depends on the value of Azimuthal quantum number. Value of m_l = +(l) to –(l).

For the “l”= 1that is “p” sub-shell, the value of “m_l” will be= +1, 0, -1.

For the “l”= 0that is “s” sub-shell, the value of “m_l” will be= 0.

For the “l”= 2that is “d” sub-shell, the value of “m_l” will be= +2, +1, 0, -1, -2.

For the “l”= 3that is “f” sub-shell, the value of “m_l” will be= +3, +2, +1, 0, -1, -2, -3.

In the above example for the outer most electron, the value of “m_l” will be +1, as it is placed in the (p_x ) orbital of the “p” sub-shell.

The orientation: Spinquantum number (m_s)

Spin quantum number

Spin quantum number tells about the orientation of the electron. Either it revolves clockwise or anticlockwise in its axis. Its value can be +1/2 or -1/2.

In the above example for the outer most electron, the value of “m_s” will be +1/2.

If you try yourself a few examples, you will find that none of the two electrons of the same atom have the similar address or in other words have the same value of all the four quantum numbers. It was first discovered by “Wolfgang Pauli”. His finding is known today as “Pauli exclusion principle”  

Correlation Table of Quantum Numbers

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Polar Covalent Bond: Water

Elements share electrons to form covalent bonds. We assume that they both have equal possession of shared pair of electrons. But it isn’t the case every time. You know each element is different to others and every element has its unique qualities, strengths and weaknesses.

Water Molecule 3D view

As you know electrons continuously move around, they look like a cloud hanging around the nucleus. When you look on the picture of water molecule, you will see the cloud is a little denser near the Oxygen atom. It means the shared electron pairs are pulled by Oxygen atom to get a greater share of it. How does Oxygen manage to do it? What makes Oxygen more powerful than Hydrogen?

The subatomic particles build an element’s strength/ weakness. You remember protons are positively charged and stay inside the nucleus. Because of these protons nucleus behaves like a magnet for electrons. Nucleus attracts electrons to bind them closer to it. Magnetic power of nucleus doesn’t work effectively on the electrons of distant orbits.

Let’s have a close look on Oxygen atom (⁸O). Its electronic configuration is 1s², 2s², 2p⁴. It has 2 orbits and 8 protons in its nucleus. So it wouldn’t be difficult for nucleus to bind electrons and even to attract electrons of bonding pair.

Oxygen Atom

When you see Hydrogen atom, it has only one proton. Its nucleus is not strong enough to withstand in the competition with Oxygen.

The capability of element to pull bonding electrons is called “electro negativity”. In water molecule Oxygen is more electro negative than Hydrogen. You don’t have to work out every time that which element is more electro negative than other. You can find this type of information about any element in the periodic table. You will be amazed to know that elements are arranged in the order of strengths and weaknesses in the periodic table. I will elaborate it in my next post.

When one element in a covalent bond is stronger than the other one, the shared electron cloud shifts towards stronger element (more electro negative) and that element develops partial negative charge. The weaker element (less electro negative) on the other hand, loses some of the cloud and develops partial positive charge. That’s how two different poles (“+” and “-“) are developed in the same molecule. Thus, covalent bond develops a bit of ionic characteristic, and such molecules are called polar molecules.

This type of polar covalent bond is the reason behind the wonderful properties of water. That’s why water is a liquid and remains liquid up to large range of temperature. It boils at quite higher temperature (100 ͦ C). Its solid form Ice is lighter than its liquid form is also due to its polar nature. It can dissolve a number of chemical substances, ions and gases. These are a few qualities of water I have mentioned. Water is the unique creation of the nature that supports life in the earth.​

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