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Why the electronic configuration of nitrogen is 1s2 2s2 2px1 2py1 2pz1?
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What if Oxygen has an atomic number of 8. The n 1 energy level has two electrons and the n 2 energy level has six electrons. Oxygen's valence is .?
Oxygen has 6 valence electrons as its electronic configuration is 2, 6.
How does radiation speed up a chemical reation?
Radiation can speed up a chemical reaction by promoting electrons to different energy levels, changing the reactivity of the atom/molecule. If radiation was not present, one may have to wait for this promotion to happen spontaneously, which can take far longer. Different electron configurations within an atom can also change the number of bonds that can be readily formed. For example, the electron configuration in carbon is 1s2 2s2 2p2. There are two unpaired electrons in the p orbital, which could readily make a total two bonds. However, carbon often makes four bonds. This occurs because one of the 2s2 electrons gets promoted to the unoccupied p orbital. The electron configuration changes to 1s2 2s1 2p3, which then has four unpaired electrons which allows it to readily form four bonds. The unpaired electrons are more obvious when the electron configuration is written in this form: 1s2 2s1 2px1 2py1 2pz1. The electrons in each of the p orbitals, as well as the 2s orbital can readily form bonds in this configuration.
Why is second electron affinity for halogens is zero?
All of the halogens are one electron short of having all of their atomic orbitals filled to reach an atom's state of nirvana. This explains why, in general, halide chemistry is such that halogens so willingly literally accept one electron in their ionic formulations and formally accept one electron or share a pair of electrons in the vast majority of their predominately covalent compounds. Halogens have no affinity for accepting a second electron because once a halogen atom has accepted once electron, all of its atomic orbitals each contain two electrons and are thus full. Any element with all its atomic orbitals filled has the equivalent electronic configuration of a noble gas and is in its most stable electronic state.What follows is very important to understand. It appears that many chemistry students do not know this fact probably because most textbooks and instructors do not explicitly point it out or they do a poor job emphasizing it: Elements only possess the atomic orbitals defined by the row in which an element exists in the Periodic Table.In many compounds, a particular element may possess one or more empty atomic orbitals in its electronic ground state. Students who have completed the first semester of general chemistry were presented with, and expected to understand, what atomic orbitals each element has. They should also know the order in which a given element's orbitals are progressively occupied by electrons when that element is in its ground electronic state and that orbitals with the lowest energy are filled first. It is also important to understand that the theoretical order of atomic orbitals in elements heavier than argon may be in a different order. This effect, when it occurs, is due to electron-electron repulsions about the element's nucleus.Let's look at a 2nd row element as an example. How about nitrogen? Because it's a 2nd row element, nitrogen has two "shells" of atomic orbitals and a total of five orbitals; however only electrons in the outer shell of orbitals may participate in chemical bonding. The 1st shell of electrons consists only of the 1s orbital. Like all atomic orbitals, the 1s orbital can hold a maximum of two electrons, which is denoted by the superscript in the orbital's designation, as in 1s2. Starting from the 1st element in the 2nd row and counting each element up to and including nitrogen shows that the outer shell of orbitals on nitrogen contains five electrons. Assuming that no electron-electron interactions alter the respective theoretical energy levels of the five orbitals (This does not occur in any of the 2nd row elements), the atomic orbitals on nitrogen are, in increasing energy: [1s2], 2s2, 2px1, 2py1, 2pz1. The three 2p orbitals have the same energy and are filled with one electron first before any of them takes on a second electron. Note that the first p orbitals, and the ones lowest in energy, are the 2p orbitals. There is simply no such thing as a 1p orbital. The 2p orbitals could have been named 1p orbitals. Everyone who first applied quantum mechanics to the hydrogen atom in order to describe its atomic emission spectrum, and, not long thereafter, the number and energy levels of an atom's electrons, are no longer with us. Nevertheless, the reason for the seemingly strange numerical designations is almost certainly because the quantum numbers that are solutions to the wave equation corresponding to the number and shape of the atomic orbitals begin with "2" for the p orbitals, "3" for the d orbitals, etc., and perhaps the people who discovered and published all of these findings decided not to change the numerical designations.The point I hope I made is that the five atomic orbitals shown for nitrogen are all it has. In addition to s and p atomic orbitals, there exists d and f orbitals, but not for nitrogen or any other second-row element. Therefore, once the 2s and 2p orbitals are filled, nitrogen cannot accept or share another additional electron because there is no atomic orbital in which it can be placed.
What is the electron configuration for N7?
1s22s22p3
How many electrons would carbon have in its outermost shell?
1 atom of the element, carbon, would have 4 electrons in its outer most shell. It's electron configuration is 1s2 2s2 2px1 2py1. 2 is the outer most shell, so there are 4 electrons in shell 2.
Ball-and stick model of a BF3 molecule?
excited state EC of boron is 1s2 2s1 2px1 2py1. so it undergoes sp2 type of hybrydisation and has trigonal planar shape.
What is the configuration for carbon atom in the molecule methane?
Electronic configuration of carbon - 1s2 2s2 2p2The valance electrons are 2s2 2p2.Actually,among 2 electron in the s orbital one is excited to p orbital,now the electronic configuration of carbon is 2s1 2px1 2py1 2pz1 .These four electrons hybridize with the incoming four hydrogen atoms (each one having 1 electron).Sp3 hybridization taking place in methane having tetrahedral geometry .
What is the Hound's rule in chemistry?
It is the rule for filling of electrons in orbitals. It states that, When orbitals of equal energy (such as px, py, pz) are available electrons, they fill orbitals singly to keep the spin parallel. For example, N(Z=7)= 1s2, 2s2, 2px2, 2py1 (wrong) N(Z=7)= 1s2, 2s2, 2px1, 2py1, 2pz1 (correct)
Electron configuration for O?
O2- would have the electron configuration of 1s2 2s2 2p6 Because the Oxygen has a charge of 2-, that means it has two more electrons than normal. So you do the electron configuration normally, then add two electrons! =]
What if Oxygen has an atomic number of 8. The n 1 energy level has two electrons and the n 2 energy level has six electrons. Oxygen's valence is .?
Oxygen has 6 valence electrons as its electronic configuration is 2, 6.
How does radiation speed up a chemical reation?
Radiation can speed up a chemical reaction by promoting electrons to different energy levels, changing the reactivity of the atom/molecule. If radiation was not present, one may have to wait for this promotion to happen spontaneously, which can take far longer. Different electron configurations within an atom can also change the number of bonds that can be readily formed. For example, the electron configuration in carbon is 1s2 2s2 2p2. There are two unpaired electrons in the p orbital, which could readily make a total two bonds. However, carbon often makes four bonds. This occurs because one of the 2s2 electrons gets promoted to the unoccupied p orbital. The electron configuration changes to 1s2 2s1 2p3, which then has four unpaired electrons which allows it to readily form four bonds. The unpaired electrons are more obvious when the electron configuration is written in this form: 1s2 2s1 2px1 2py1 2pz1. The electrons in each of the p orbitals, as well as the 2s orbital can readily form bonds in this configuration.
Find an electron dot diagram for nA2O?
i think it's meant to be NO4 3- .. an ion because it's NO2 - and NO3 2- but then i don't know how they're actually bonded! i guess there's deffo one dative bond between N and O and then either 3 dot cross ones ..or another dative from n to o and one from an o to n and one dot cross! but i expect it's three dot crosses and one dative! because it's more even.. probs. guut luck anyway :D i think it's meant to be NO4 3- .. an ion because it's NO2 - and NO3 2- but then i don't know how they're actually bonded! i guess there's deffo one dative bond between N and O and then either 3 dot cross ones ..or another dative from n to o and one from an o to n and one dot cross! but i expect it's three dot crosses and one dative! because it's more even.. probs. guut luck anyway :D
Why is second electron affinity for halogens is zero?
All of the halogens are one electron short of having all of their atomic orbitals filled to reach an atom's state of nirvana. This explains why, in general, halide chemistry is such that halogens so willingly literally accept one electron in their ionic formulations and formally accept one electron or share a pair of electrons in the vast majority of their predominately covalent compounds. Halogens have no affinity for accepting a second electron because once a halogen atom has accepted once electron, all of its atomic orbitals each contain two electrons and are thus full. Any element with all its atomic orbitals filled has the equivalent electronic configuration of a noble gas and is in its most stable electronic state.What follows is very important to understand. It appears that many chemistry students do not know this fact probably because most textbooks and instructors do not explicitly point it out or they do a poor job emphasizing it: Elements only possess the atomic orbitals defined by the row in which an element exists in the Periodic Table.In many compounds, a particular element may possess one or more empty atomic orbitals in its electronic ground state. Students who have completed the first semester of general chemistry were presented with, and expected to understand, what atomic orbitals each element has. They should also know the order in which a given element's orbitals are progressively occupied by electrons when that element is in its ground electronic state and that orbitals with the lowest energy are filled first. It is also important to understand that the theoretical order of atomic orbitals in elements heavier than argon may be in a different order. This effect, when it occurs, is due to electron-electron repulsions about the element's nucleus.Let's look at a 2nd row element as an example. How about nitrogen? Because it's a 2nd row element, nitrogen has two "shells" of atomic orbitals and a total of five orbitals; however only electrons in the outer shell of orbitals may participate in chemical bonding. The 1st shell of electrons consists only of the 1s orbital. Like all atomic orbitals, the 1s orbital can hold a maximum of two electrons, which is denoted by the superscript in the orbital's designation, as in 1s2. Starting from the 1st element in the 2nd row and counting each element up to and including nitrogen shows that the outer shell of orbitals on nitrogen contains five electrons. Assuming that no electron-electron interactions alter the respective theoretical energy levels of the five orbitals (This does not occur in any of the 2nd row elements), the atomic orbitals on nitrogen are, in increasing energy: [1s2], 2s2, 2px1, 2py1, 2pz1. The three 2p orbitals have the same energy and are filled with one electron first before any of them takes on a second electron. Note that the first p orbitals, and the ones lowest in energy, are the 2p orbitals. There is simply no such thing as a 1p orbital. The 2p orbitals could have been named 1p orbitals. Everyone who first applied quantum mechanics to the hydrogen atom in order to describe its atomic emission spectrum, and, not long thereafter, the number and energy levels of an atom's electrons, are no longer with us. Nevertheless, the reason for the seemingly strange numerical designations is almost certainly because the quantum numbers that are solutions to the wave equation corresponding to the number and shape of the atomic orbitals begin with "2" for the p orbitals, "3" for the d orbitals, etc., and perhaps the people who discovered and published all of these findings decided not to change the numerical designations.The point I hope I made is that the five atomic orbitals shown for nitrogen are all it has. In addition to s and p atomic orbitals, there exists d and f orbitals, but not for nitrogen or any other second-row element. Therefore, once the 2s and 2p orbitals are filled, nitrogen cannot accept or share another additional electron because there is no atomic orbital in which it can be placed.
What type of bond is formed when you combine chlorine and carbon?
When carbon(C) combines with chlorine(Cl), carbon forms single covalent bond with each chlorine atom. The valence shell configuration of uncombined C atom is 2s2 2px1 2py1 .During combination hybridisation takes place, an electron shifts from 2s to empty 2pz . These four orbitals merge to form four hybrid orbitals. Each hybrid orbital has 1 electron which pairs up with a valence electron of Cl atom to form a covalent bond. Thus the four hybrid orbitals form four single covalent bonds with four chlorine atoms. This completes each atom's octet and gives them noble gas configuration.When two atoms of different elements combine they always have electronegativity(EN) difference due to which the shared pair of electrons is more closer to more electronegative atom. This gives the bonds some ionic character. So no bond between two atoms of different elements can be purely covalent.Percentage of ionic character is found by following formula:16*|χA - χB| + 3.5*(χA- χB)2where χA and χB are electronegativity of the two combining atoms.For C-Cl bond percentage of ionic character is 8.87 (approx).
Shorthand orbital notation?
1s2 2s2 2p4 Meaning 2 electrons in the 1s orbital/shell, 2 electrons in the 2s orbital & 4 electrons in the 2p. So both 1s & 2s orbitals are full, the 2p orbital is only partly filled as it can hold 6 electrons.
