Hydrogen electron configuration is 1s1. The electron configuration of hydrogen shows that hydrogen is an s-block element. The valency and valence electrons of hydrogen are 1.
This article gives an idea about the electron configuration of hydrogen, the period and groups of hydrogen, the valency of hydrogen, and valence electrons, hydrogen bond formation, compound formation, application of different principles.
The first element of the periodic table is hydrogen. And its position at the beginning of the periodic table. The periodic table begins with hydrogen. The electron configuration of each atom is done in 2 ways. Therefore, the hydrogen electron configuration can be done in two ways.
1. Electron configuration through orbit.
2. Electron configuration through orbital.
Electron configuration via orbitals follows different principles. For example, the Aufbau principle, Hund’s principle, Pauli’s exclusion principle.
Hydrogen Electron Configuration through orbit
Scientist Niels Bohr provided a model of the atom in 1913. The complete idea of the orbit is given there. The electrons of the atom revolve around the nucleus in a certain circular path. These circular paths are called orbit. These orbits are expressed by n. [n = 1,2 3 4. . .]
K is the name of the first orbit, L is the second, M is the third, N is the name of the fourth orbit. The electron holding capacity of each orbit is 2n2. [Where, n = 1,2 3,4 . . .]
n = 1 for K orbit.
The electron holding capacity of K orbit is 2n2 = 2×12 = 2 electrons.
For L orbit, n = 2.
The electron holding capacity of the L orbit is 2n2 = 2×22 = 8 electrons.
The atomic number of hydrogen is 1. That is, the number of electrons in hydrogen is one. We know that the first orbit has two electrons holding capacity but the total number of electrons in hydrogen is 1. Therefore, 1 electron exists in the first orbit of hydrogen. And 1 electron is located in the first orbit.
Hydrogen electron configuration in the Aufbau principle
Hydrogen electron configuration in the Aufbau principle is the arrangement of electrons through orbitals. The Aufbau principle is that the electrons present in the atom will first fill the low energy orbital and then gradually fill the high energy orbital. These orbitals are – s, p, d, f.
The electron holding capacity of these orbitals is s = 2, p = 6, d = 10, f = 14. The Aufbau electron configuration method is, 1 s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d. The electron configuration of hydrogen in the Aufbau method is 1s1.
Hund’s principle for hydrogen electron configuration
Hund’s principle is that when electrons enter the orbitals of equal power, the electrons will enter the orbital in a random manner as long as the orbital is empty. And the spin of these unpaired electrons will be one-sided. But, s sub-orbit has only one orbital. That is why Hund’s principle does not support the entry of new electrons into the s-orbital.
Hydrogen electron configuration is 1s1. The last orbital of hydrogen has one unpaired electron but the number of orbitals is only one. Therefore, the hydrogen atom does not support Hund’s principle.
Determining the period and group of hydrogen by electron configuration
The last orbital of an element is the period of that element. The electron configuration of hydrogen shows that the last orbital of hydrogen is 1(1s). That is, the period of hydrogen is 1.
Again, the total number of electrons in the last orbit of an element is the group of that element. The total number of electrons in the last orbit of hydrogen is 1. Therefore, the group number of hydrogen is 1. Therefore, the period and group of hydrogen are both 1.
Determining the valency and valence electrons of hydrogen
The hydrogen electron configuration is 1s1. If the last orbit of an element has 1,2,3 or 4 electrons, then the number of electrons in the last orbit is the valency of that element. From the hydrogen electron configuration, we can say that 1 electron exists in the last orbit of hydrogen.
Therefore, the valency of hydrogen is 1. Again, the number of electrons in the last orbit of an element, the number of those electrons is the valence electrons of that element.
In the electron configuration of hydrogen, we see that 1 electron exists in the last orbit of hydrogen. Therefore, the valence electrons of hydrogen are 1. Finally, we can say that the valency and valence electrons of hydrogen are 1.
Determining the block of hydrogen by electron configuration
The elements that have the last electron entering the s orbital after electron configuration are called s-block elements. Again, the elements in group 1 of the periodic table are the s-block elements.
The hydrogen electron configuration is 1s1. The electron configuration of hydrogen implies that the last electron of hydrogen enters the s orbital (1s) . As we know, hydrogen is an element of group 1 and the last electron of hydrogen enters the s orbital. So, we can say that hydrogen is the s block element.
Formation of hydrogen ionic bonds
The bond formed by the creation of positive and negative ions is called an ionic bond. During the chemical connection of metal and non-metal atoms, one or more electrons of the last energy level of the metal atom are transferred to the last energy level of the non-metal atom.
The compound that forms ionic bonds is called an ionic compound. Magnesium atoms form ionic compounds with hydrogen. The electron configuration of magnesium(Mg) is 1s 2 2s 2 2p 6 3s 2.
The electron configuration of the magnesium atom shows that there are 2 electrons in the last orbit of magnesium. The magnesium atom acquires a stable octal structure of near inert gas by releasing 2 electrons from its last orbit. And magnesium is converted to Mg+2 ions.
On the other hand, 1 electron exists in the last orbit of the hydrogen atom. The hydrogen atom acquires the structure of helium by accepting 1 electron. And hydrogen is converted to H– ions. Inversely charged Mg2+ and 2H– ions are combined by attraction to each other to form an MgH2 compound through ionic bonding.
Covalent bonding of hydrogen
The bond formed by the electron share between two atoms is called a covalent bond. The hydrogen atom combines with the carbon, fluorine, chlorine, oxygen, and silicon atoms to form covalent bonds. And (CH4, HF, HCl, H2O) form compounds.
In the case of H2O: The electron configuration of hydrogen and oxygen atoms is-
The hydrogen electron configuration is 1s1.
And oxygen electron configuration is 1s2 2s2 2p4.
The electron configuration of hydrogen shows that 1 electron exists in the hydrogen atom. Again, the electron configuration of oxygen shows that there are 6 electrons in the last orbital of the oxygen atom. Two hydrogen atoms join one oxygen atom to form a covalent bond through electron sharing. And forms H2O compounds through covalent bonds.
In the case of SiH4: The electron configuration of silicon(Si) and hydrogen(H) is-
The electron configuration of silicon is- Si(14) = 1s2 2s2 2p6 3s2 3p2.
The electron configuration of hydrogen is- H (1) = 1s1.
There are 4 electrons in the last orbit of a silicon atom. These four electrons in the silicon atom cannot be abandoned or accepted. So, a silicon atom shares four electrons with four hydrogen atoms. Si-H forms single covalent bonds and SiH4 forms compounds.
Properties of hydrogen
- The atomic number of hydrogen is 1.
- The total number of electrons in hydrogen is 1.
- The active atomic mass of hydrogen is [1.00784, 1.00811].
- Hydrogen is an s-block element.
- s-block elements react with hydrogen to form hydride compounds.
- The valency of hydrogen is 1.
- Hydrogen compounds are highly alkaline.
- The valency and valence electrons of hydrogen are 1.
- The group and period of hydrogen are both the same.
- Hydrogen forms ionic and covalent bonds.
- The melting point of hydrogen is 13.99 K (−259.16 °C, −434.49 °F) and the boiling point is 20.271 K (−252.879 °C, −423.182 °F).
- The electronegativity of hydrogen is 2.20
Reaction with hydrogen
Hydrogen reaction with group-1 elements
The elements of group-1 are Lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs). Which is known as an alkali metal. Alkali metals react with dry hydrogen at 400 ° C to form metallic hydride compounds. Li, Na, K, Cs all the elements of group-1 react with hydrogen to form hydride compounds.
2Na (s) + H2 (g) → 2NaH ( Na+ + H– ) (sodium hydride)
2K (s) + H2 (g) → 2KH (potassium hydride)
2Rb (s) + H2 (g) → 2RbH (Rubidium hydride)
2Cs + H2 → 2CsH (cesium hydride)
But in the case of the Li atom, its temperature is 800 ° C.
2Li + H2 → 2LiH (lithium hydride).
Going from top to bottom(Li to Cs), the activity of the reaction decreases. This is because the cation size increases from the top to the bottom of the group. As a result, the M – H bond is weakened. [ Here, M = Li, Na, K, Cs, Rb ]
Hydrogen reaction with Group-2 elements
Group-2 elements are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). All the elements except the beryllium (Be) element of group-2 react with hydrogen. And forms hydride compounds.
Mg (s) + H2 → MgH2
Ca (s) + H2 → CaH2
Sr (s) + H2 → SrH2
The reaction of halogen with hydrogen
The elements in group-17 are halogens. The halogen elements are fluorine (F), chlorine (Cl), bromine (Br), iodine (I), astatine (At). Each halogen atom reacts with hydrogen and forms compounds.
H2 + F2 → 2HF
H2 + Cl2 → 2HCl
H2 + Br2 → 2HBr
H2 + I2 → 2HI
All the above compounds are soluble in water. The above compounds (H–Cl, H–F, H–Br, H–I) are mixed with water to form H+ ions. Which combines with water to produce H3O + ions.
HF + H2O → H3O + + F –
HCl + H2O → H3O + + Cl –
HBr + H2O → H3O + + Br –
HCl + H2O → H3O + + I –
Formation of hydride compounds
Under special conditions, at high pressures and temperatures, the elements of group-15 combine with hydrogen to form hydride compounds. E.g., NH3, PH3, AsH3.
N2 + 3H2 → 2NH3 + heat. (200atm and 500°C)
H2 does not react directly with phosphorus. However, PH3 is produced when white phosphorus is heated together in a solution of caustic soda.
P4 (white) + 3NaOH + 3H2O → PH3 + 3NaH2PO2
P4 (white) + 3KOH + 3H2O → PH3 + 3KH2PO2
The two compounds are Lewis base because of the presence of unpair electrons in the NH3 and PH3 molecules. The alkalinity of NH3 is higher than that of PH3.
The value of electronegativity of N atom in NH3 molecule is 3.0 but the value of electronegativity of P atom in PH3 atom is 2.1. That is, the Nitrogen(N) atom is more negatively charged than the P atom.
Therefore, the concentration of electrons in the N–H bond between the NH3 molecules tends to be more towards the N atom. NH3 compounds are more alkaline than PH3.
Application of Hydrogen Dual Law
Gaining the electron configuration of the last orbital of atoms through the exchange and sharing of electrons between the elements themselves is called dual law.
In the case of H2: Hydrogen electron configuration shows that 1 electron of hydrogen exists. The hydrogen atom receives one electron and acquires the electron configuration of helium and becomes more stable by exhibiting the same properties as an inert gas. The two hydrogen atoms form the H2 compound through electron share.
Explanation of hydrogen bonding
Hydrogen bond: When a hydrogen atom combines with a very high electrically negative element to form a covalent compound, the electrons participating in the bond are more attracted to a very high electronegative element.
As a result, polarity is created between them. When such polar molecules come close to each other, the positive hydrogen end is particularly attracted to the negative end of the other molecule and forms a bond through a weak attraction. This weak attraction is called hydrogen bonding.
The hydrogen bond is expressed by the dot (. . .) sign. A hydrogen bonding strength is about 0.01. For example, Hydrogen bonding is observed in molecules like hydrogen fluoride (HF), water (H2O), ammonia (NH3), ethanoic acid (CH3COOH), phenol (C6H5OH), etc.
Properties of hydrogen bond
- Hydrogen bonds are weak bonds. Even more so than covalent bonds. Hydrogen bonding strength 8–42 kJ/mol. On the other hand, the covalent bond strength is 200–450kj/mol.
- The strength of the hydrogen bond depends on the value of the electronegativity of the hydrogen atom to the other atom. The higher the value of electronegativity of the connected atom, the higher the hydrogen bonding force. The values of electronegativity of elements F, O, and N are 4.0, 3.5, and 3.0. A sequence of hydrogen bonding energy H — F> H — O> H — N.
- Hydrogen bonds have specific bonding orientations.
- The position of the hydrogen bond depends on the direction of the unpair electron present in the atom of the electronegative element associated with the hydrogen atom.
- A large number of molecules are joined together by hydrogen bonding. As a result, the molecules remain in a cohesive state.
- The physical position of the molecule changes in hydrogen bonding.
- The size of a hydrogen atom is smaller than the size of an atom of another electronegative element bound by a hydrogen bond. As a result, a strong repulsive force acts between the last orbital electrons of the two atoms.
- In order to minimize the value of repulsion, the positions of hydrogen and electronegativity are linear.
- The effect of hydrogen bonding is to change the melting point, boiling point, solubility, density, viscosity, the surface texture of the compound.
Prerequisites for Hydrogen Bonding
- The corresponding molecules must have hydrogen atoms.
- The atom attached to the hydrogen atom in the corresponding molecule must be an extremely electronegative element. E.g., O, F, N.
- Molecules must have unpairs electrons.
- The influence of unpaired electrons plays a major role in the formation of hydrogen bonds.
- The bond between the electronegative element and the hydrogen atom must be more polarized.
- The size of the electronegative atom attached to the hydrogen atom must be small. The smaller the size of the atom attached to the hydrogen atom, the more effective the polarization between the positive edge of the hydrogen atom and the negative edge of the electronegative element becomes more effective. The efficiency of hydrogen bonding is also increased. For this reason, Cl, Br, S, and P form compounds with negative elements hydrogen but do not form hydrogen bonds.
- Must have static electron attraction. This attraction results in the formation of hydrogen bonds.
Types of Hydrogen Bonding
There are two types of hydrogen bonds.
i. Intermolecular hydrogen bonding.
ii. Intramolecular hydrogen bonding.
Intermolecular hydrogen bonding
Hydrogen bonds formed between different molecules of the same or different compounds are called intermolecular hydrogen bonds. Hydrogen bonds are formed between individual molecules of the same or different compounds. E.g. HF, H2O, CH3COOH.
Intramolecular hydrogen bonding
Hydrogen bonding between different parts of the same molecule of the same compound is called intramolecular hydrogen bonding. The formation of hydrogen bonds is called chilation. E.g., C6H4(OH)(NO2), hydroxy-benzaldehyde C6H4(OH)CHO, salicylic acid C6H4(OH)COOH. Intramolecular hydrogen bonds exist between these molecules.
The main topic of this article is the hydrogen electron configuration. Hydrogen is the first element of the periodic table. The hydrogen element is exceptional compared to the other elements in the periodic table.
This article gives an idea about the electron configuration of hydrogen, the stages and groups of hydrogen, hydrogen valence, and valence electron, hydrogen bond formation, hydrogen compound formation, application of different principles.
1. How do you write the electron configuration for hydrogen?
Ans: Hydrogen electron configuration is 1s1.
2. What is the electron configuration for the H+ ion?
Ans: Electron configuration for the H+ ion is 1s0.
3. What is the formula of hydrogen(H+) ion?
4. How many electron does H+ have?
Ans: 0 electron.
5. Is H+ just a proton?
Ans: yes. H+ is just a proton.
6. How many valence electrons does hydrogen have?
Ans: 1 valence electron.
7. How many valence electrons does H+ have?
Ans: zero valence electron.
- Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. p. 240. ISBN 978-0123526519.
- Lide, D. R., ed. (2005). “Magnetic susceptibility of the elements and inorganic compounds”. CRC Handbook of Chemistry and Physics (PDF) (86th ed.). Boca Raton (FL): CRC Press. ISBN 978-0-8493-0486-6.