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Topic: **Help me with my chemistry homework****Question:**
OK, so I have this chemistry homework and I have no Idea how to do it. I have tried so hard to get it but i cant, does anybody know how to do electronic configurations and the quantum number stuff. But I mostly just need help with the configuration part.
if you can help me i would appreciate it soo much.. thanks
my question was how do you do electron configurations??

June 16, 2019 / By Susana

Electronic configurations have to do with orbitals and electrons. Any atom starts off with a 1s orbital, and every additional shell can have an extra orbital. So the second shell can have two orbitals, s and p, and they are labelled 2s and 2p (2 for 2nd shell). The third shell can have 3 (3s, 3p and 3d), but there's a catch. The 4s orbital is lower in energy than the 3d, so the 4s is filled first (don't know why). Each additional orbital contains two more sub-orbitals (I don't know if that's the correct word) than the previous, and can contain two electrons in each. This means that every s-orbital can have 2 electrons, but a p-orbital (consisting of 3 lobes) can have 3 × 2 = 6 electrons, and a d-orbital can have 5 × 2 = 10 electrons, and so on. As far as quantum numbers are concerned, each electron has 4. The principle quantum number, n, can go up to the number of shells in the atom, and will be equal to or lower than the period number of the atom. The orbital quantum number, l, designates the type of orbital. It will always be lower than n, and never lower than zero. l = 0 means the s-orbital, l = 1 means p, etc. The magnetic quantum number, m(l) (l is subscripted) means the sub-orbital. Its value can range between the negative and positive values of l. E.g, if l = 2, m(l) can be -2, -1, 0, 1 or 2. Finally the spin quantum number, m(s), is either 1/2 or -1/2 for an electron. For example, a neutral atom of nitrogen has electron configuration 1s2 2s2 2p3, whereas manganese will have 1s2 2s2 2p6 3s2 3p6 4s2 3d5. Hope this helps a bit.

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Electronic configurations have to do with orbitals and electrons. Any atom starts off with a 1s orbital, and every additional shell can have an extra orbital. So the second shell can have two orbitals, s and p, and they are labelled 2s and 2p (2 for 2nd shell). The third shell can have 3 (3s, 3p and 3d), but there's a catch. The 4s orbital is lower in energy than the 3d, so the 4s is filled first (don't know why). Each additional orbital contains two more sub-orbitals (I don't know if that's the correct word) than the previous, and can contain two electrons in each. This means that every s-orbital can have 2 electrons, but a p-orbital (consisting of 3 lobes) can have 3 × 2 = 6 electrons, and a d-orbital can have 5 × 2 = 10 electrons, and so on. As far as quantum numbers are concerned, each electron has 4. The principle quantum number, n, can go up to the number of shells in the atom, and will be equal to or lower than the period number of the atom. The orbital quantum number, l, designates the type of orbital. It will always be lower than n, and never lower than zero. l = 0 means the s-orbital, l = 1 means p, etc. The magnetic quantum number, m(l) (l is subscripted) means the sub-orbital. Its value can range between the negative and positive values of l. E.g, if l = 2, m(l) can be -2, -1, 0, 1 or 2. Finally the spin quantum number, m(s), is either 1/2 or -1/2 for an electron. For example, a neutral atom of nitrogen has electron configuration 1s2 2s2 2p3, whereas manganese will have 1s2 2s2 2p6 3s2 3p6 4s2 3d5. Hope this helps a bit.

basically the easiest way i can explain it is to work across the periodic table. start with ns1 and ns2 for the first two elements in each row, then in rows 2 and 3 you move to np1 through np6. when you get to rows 4 and 5, you start with the ns2, but then you move to nd10. **nd10 the n will be the row before the one you're on, i'll put an example at the end. then put the np6 at the end. when you get to rows 6 and 7 you start with ns2, but put nf14 before nd10. the n for the f sublevels are 2 less than the n that is in the ns2. then you put np6 again. for all of them, ns2 and np6 - n is the row you're on. ex: lead: Pb is 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2. i stopped with 6p2 because that's where Pb is on the periodic table, it's not all the way on the right with the noble gases. hope this helps!!

basically the easiest way i can explain it is to work across the periodic table. start with ns1 and ns2 for the first two elements in each row, then in rows 2 and 3 you move to np1 through np6. when you get to rows 4 and 5, you start with the ns2, but then you move to nd10. **nd10 the n will be the row before the one you're on, i'll put an example at the end. then put the np6 at the end. when you get to rows 6 and 7 you start with ns2, but put nf14 before nd10. the n for the f sublevels are 2 less than the n that is in the ns2. then you put np6 again. for all of them, ns2 and np6 - n is the row you're on. ex: lead: Pb is 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2. i stopped with 6p2 because that's where Pb is on the periodic table, it's not all the way on the right with the noble gases. hope this helps!!

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Electronic configurations have to do with orbitals and electrons. Any atom starts off with a 1s orbital, and every additional shell can have an extra orbital. So the second shell can have two orbitals, s and p, and they are labelled 2s and 2p (2 for 2nd shell). The third shell can have 3 (3s, 3p and 3d), but there's a catch. The 4s orbital is lower in energy than the 3d, so the 4s is filled first (don't know why). Each additional orbital contains two more sub-orbitals (I don't know if that's the correct word) than the previous, and can contain two electrons in each. This means that every s-orbital can have 2 electrons, but a p-orbital (consisting of 3 lobes) can have 3 × 2 = 6 electrons, and a d-orbital can have 5 × 2 = 10 electrons, and so on. As far as quantum numbers are concerned, each electron has 4. The principle quantum number, n, can go up to the number of shells in the atom, and will be equal to or lower than the period number of the atom. The orbital quantum number, l, designates the type of orbital. It will always be lower than n, and never lower than zero. l = 0 means the s-orbital, l = 1 means p, etc. The magnetic quantum number, m(l) (l is subscripted) means the sub-orbital. Its value can range between the negative and positive values of l. E.g, if l = 2, m(l) can be -2, -1, 0, 1 or 2. Finally the spin quantum number, m(s), is either 1/2 or -1/2 for an electron. For example, a neutral atom of nitrogen has electron configuration 1s2 2s2 2p3, whereas manganese will have 1s2 2s2 2p6 3s2 3p6 4s2 3d5. Hope this helps a bit.

basically the easiest way i can explain it is to work across the periodic table. start with ns1 and ns2 for the first two elements in each row, then in rows 2 and 3 you move to np1 through np6. when you get to rows 4 and 5, you start with the ns2, but then you move to nd10. **nd10 the n will be the row before the one you're on, i'll put an example at the end. then put the np6 at the end. when you get to rows 6 and 7 you start with ns2, but put nf14 before nd10. the n for the f sublevels are 2 less than the n that is in the ns2. then you put np6 again. for all of them, ns2 and np6 - n is the row you're on. ex: lead: Pb is 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2. i stopped with 6p2 because that's where Pb is on the periodic table, it's not all the way on the right with the noble gases. hope this helps!!

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