Valence Electrons and Valency of an Element

Valence electrons play an important role in an atom’s life. When an atom wants to make a bond with another atom, these Valence electrons make it possible. Valence electrons are placed in the outer most orbit of the atom and that orbit (shell) is called the valence shell. Let’s take an example to identify them. Write the configuration of Oxygen

⁸O : 1s² , 2s² , 2p⁴

Oxygen has 6 electrons in its outer most orbit (2+4) that means it has 6 valence electrons and orbit 2 is called the valence shell. Now place these electrons in orbitals of their respective sub-shell.

You can see, here 2 sets of paired electron and 2 single electrons are present. Paired electrons are called the lone pair. In Lewis dot structure these lone pairs are drawn as paired dots while other valance electrons as single dots.

Oxygen has 6 valence electrons and it needs 2 more electrons to achieve octet. The number of electrons needed to complete octet is termed as the valency of an element. This value can be obtained by subtracting the number of valence electrons from octet (8). So, the valency of Oxygen is 2. That means Oxygen is ready to share its 2 valence electrons or ready to accept 2 more electrons in order to achieve octet.

To grasp the concept of valence electrons and valencies of elements lets work out a few examples. Before working on examples, I want to tell you one more thing about periodic table. Valence electrons and valencies are also periodic properties. All elements of a group have same number of valence electrons and valency. And you can call it as group valency. Let’s figure out it for the first element of each group.

Every element has a right to choose which way it likes to achieve octet. It has always two choices either it can accept electrons from others or donate to others. It chooses between them as per its financial condition (in case of atom, it means energy required). If donation of electrons requires lesser energy (lower IE) it chooses to donate and if it can manage to take loan from others (lower EGE) it chooses to accept electrons. 

Elements of group 1 have 1 electron in outer most orbit. They have two ways to achieve octet, either they can accept 7 electrons or donate 1 electron to another element. If you recall the trend of Ionisation enthalpy in periodic table, you will find that group 1 has lowest IE. That’s why they choose to donate 1 electron

and the valency of first group is 1. And for similar reason group 2 and 13 also choose to donate electrons and their respective valencies are 2 and 3.

Elements of group 14 have 4 electrons in outer orbit. For this group IE and EGE both have moderate values. That’s why they can happily donate or accept electrons from others. This quality makes them special and because of this speciality carbon has capability to form bond with a large number of elements and is able to form a number of compounds. The list of carbon compounds is so long that scientists have to develop a whole branch of chemistry to study them. This branch is known as Organic Chemistry.

Elements of group 15, 16 and 17 have higher IE but lower EGE, that’s why they prefer to accept electrons from other elements to achieve octet. And their respective group valencies are 3, 2 and 1.

Elements of group 18 have complete octet that’s why they have 0 valencies.

Do you notice that Group 1 prefers to donate 1 electron while group 17 prefers to accept 1 electron from others?  And group 1 and group 17 both have group valencies of 1. How do you differentiate them?

Scientist also faced this problem and they developed a method to clear this confusion. They give plus (+) sign to the valency when electron is donated by the element and they give minus (-) sign to the valency when electron is accepted by the element. Now valency of group 1 gets ‘+’ sign while valence of group 17 gets ‘-’ sign. This new look of valency is known as “Oxidation Number”

Do you know that periodic table tells us which element is likely to bond with which element and also gives us a clue about the nature of bond? In the next post we will explore this quality of periodic table.

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Periodic Properties at a Glance

We have discussed all periodic properties like size/ atomic radius, ionization enthalpy, electron gain enthalpy and electronegativity separately. All of them are interrelated to each other. In this post we will try to find relation between all periodic properties.

The secret of every property of an element lies within its atom. So you must have to consider the structure of an atom and arrangement of its subatomic particles.

When you go towards right in a period, number of orbits remains the same but the increasing number of protons makes nucleus stronger, therefore the size of atom decreases, as a result of which EN increases. As you know, it is easier to add an extra electron to a more electronegative element, hence the EGE decreases. And since ionisation is just opposite to EGE, the IE increases on going towards right in a period.

When you go downwards in a group, number of orbits increases, hence the size of atom increases which diminishes the power of nucleus and that’s why it becomes easier to remove an outermost electron from the neutral atom that means IE decreases. Since electron gain is opposite to the ionization so, EGE increases downward. And because of increased size it becomes difficult for nucleus to attract bonding electrons that means EN decreases downwards in a group.

You have seen that how subatomic particles play an imp role in deciding the nature of an atom, which in turn decides the nature of an element. Electrons are the most important part of an atoms life. They decide its stability and help it to achieve octet. The electrons that help in bonding are special electrons. These special electrons are called Valence electrons. In the next post we will learn more about Valence electrons.

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How Does Electronegativity Decide The Nature of Bond?

I hope now you have developed an understanding about atom and its behaviour. Just like we want financial stability in our life, atoms want stable electronic configuration. Every atom wants to achieve octet (8 electrons in outermost orbit). They try their best to achieve octet, either they borrow/ donate or share electrons. Do you know the best thing about atoms? They never achieve octet alone. They always achieve octet in pairs. They help each other to achieve octet by making bonds.

If one poor atom has 7 electrons and needs one more electron to achieve octet, then it will look for a partner who is richer than it,  has one extra electron and also willing to donate one electron to achieve octet. In this situation they bond together by gain and loss of electron so that both of them can achieve octet. Here one atom gives its electrons to other. This type of bond is called Ionic Bond.

If one poor atom has 7 electrons and needs one more electron to achieve octet but it can’t find a richer atom that has extra electron, then it will look for a companion who is similar to it and have similar needs. Both of these poor companions make a bond together by sharing their electrons and helping each other to achieve octet. This type of bond is called Covalent Bond.

We have learnt that there are two kinds of bond. One is “Ionic Bond” which we have studied in case of NaCl the table salt and other is “Covalent Bond” which we have studied in case of water molecule H2O.

Do you identify this poor atom? This poor atom is Chlorine ¹⁷Cl.

¹⁷Cl: 1s², 2s², 2p⁶, 3s², 3p⁵

It has 7 electrons in the outermost orbit. It needs to succeed in finding a richer atom willing to donate one electon. (guess who?) Sodium ¹¹Na.

¹¹Na: 1s², 2s², 2p⁶, 3s¹

It has 8 electrons in second last (penultimate) orbit and 1 electron in last orbit. And if you recall the trend of Ionization Enthalpy you will find Na belongs to group 1 and group 1 has the least IE. That means it is easiest to remove an electron from Na.

And if you recall the trend of Electron Gain Enthalpy (Electron affinity) you will find that Cl belongs to 17^th group which has larger negative values of EGE. That means it will accept one more electron quite happily and will release large amount of energy.

Now both Cl and Na have found themselves suitable for each other to fulfil their respective needs. The bond is formed between them by loss of 1 e^- from Na and gain of 1 e^- by Cl. When you compare Electronegativity values of them you will find Cl is more electronegative than Na. So Cl will Keep all 100% share of bonding electrons.

In other case if Cl couldn’t find a richer partner who is willing to donate an electron, it finds another Cl and both share their one valence electron and make a bond. This bond is made by sharing of electrons and both Cl atoms have 50% share of bonding electron pair.

When a covalent bond is formed between two different elements, bonding electrons are shared between them but it isn’t necessary that these are shared by them equally. Bonding electrons are obliged to revolve around the orbits of the both bonded atoms. More electronegative element attracts bonding electrons and compel them to spend more time in its orbit, therefore that element develops small negative charge (delta negative δ^-) and consequently the other element will get small positive charge (delta positive δ⁺). I am emphasising here on small charge because the element gets partial possession on bonding electrons; if it gets full possession, it will develop whole negative charge. By developing δ^- and δ⁺charges, a covalent bond develops a polar character.

We have seen two extreme situations above, in ionic bond when one of the bonded atoms keeps all 100% share and in covalent bond when both bonded atoms have got 50% share. The percentage share of bonding electrons between bonded atoms decides the nature of bond.

If bonded elements have similar electronegativity, that is equal tendency to attract bonding electrons, the bond between them will be fairly covalent in nature, which means both elements have 50% share of bonding electron. If bonded elements have electronegativity difference more than 1.7 in Pouling scale, the more electronegative element gets the 100% share of bonding electrons and the bond will be purely ionic in nature. If bonded elements have electronegativity difference less than 1.7 in Pouling scale, the more electronegative element gets the more than 50% share of bonding electrons and the bond will be partly ionic and partly covalent in nature, which means bond has developed polar character, like in case of water molecule.

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Periodic Property: Electro Negativity

We have learnt about the atom and discussed different qualities of elements like size, ionization enthalpy and electron gain enthalpy. Up till now we have studied an isolated element as it exists alone. When an element makes bond with other element, its other qualities are surfaced and it behaves differently in response to the changed environment. 

Like when a child interacts with the other children for the first time, his personality is explored, either he dominate others or is dominated by others or may become a believer of equality. Similarly when an element makes a bond with other element, its hidden qualities appear. Electro negativity is one of such qualities.

When two atoms make bond with each other they share bonding electrons but there is always a fight over getting a greater share of bonding electrons. Each of bonded atoms tries harder to pull bonding electrons towards it to get the major share. It is the game of strength and whichever has the greater strength to pull bonding electrons, wins and get the greater share of it. This strength is called the Electronegativity.

Electronegativity (EN) is the tendency of an atom to pull bonding electrons toward itself. As in case of water molecule, Oxygen atom pulls bonding electrons.

In 1930, scientist Pauling defined electronegativity of an atom as the tendency of the atom to attract electrons to itself when combined in a compound. This quality of element also depends on the atom and its characteristics.

• Size of the atom: In a smaller atom orbits are placed closely to the nucleus as compared to a larger one. So the nucleus is able to attracts bonding electrons effectively.
• Nuclear charge: Large nuclear charge can attract bonding electrons effectively.
• Shielding effect: 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 while, weak shielding enables nucleus to attract outer electrons as well as bonding electrons more powerfully.
• Electronic configuration: Half filled and fulfilled sub-shells have extra stability. Elements those are nearing their half filled or full filled configuration have strong desire to get possession on bonding electrons so that they can achieve stable configuration.

When you go downwards in a column, the number of protons increases which make nucleus stronger, but atomic size also increases, which diminishes the power of nucleus. So, it becomes difficult for nucleus to attract bonding electrons. You will find elements on the bottom of a column are less electronegative as compare to the elements on the top.

When you go towards right in a row, size/ radius of the atoms decreases. Increasing number of protons makes nucleus stronger. As the number of orbits remains same in a row, it becomes easier for nucleus to attract bonding electrons. That’s why; elements on the right end of a row are more electronegative than the elements on the left end.

Electronegativity is a relative quality that’s why it has no unit. The electronegativity values we are referring today’s are derived by Pouling. According to the Pouling scale group 1 has lowest EN value and group 17 has highest EN value.

EN plays important role in deciding the nature of bond. In next post will see how EN decides the nature of bond.

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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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