Can someone check for me if this is correct? I got a few wrong on my first try, so here's the new one.

cesium- 6s^1
rutherfordium- 7s^2 6d^2
cadmium- 5s^2 4d^10
silicon- 3s^2 3p^2
ununhexium 7s^2 7d^10 7p^4
lawrencium 7s^2 5f^14 6d^10 7p^4
xenon 5s^2 4d^10 5o^6

thanks!

for cadmium i think u left out the p

This site will be your friend.

Click "Orbitals" from the top menu and hover over an element. It has a bunch of stuff that will help for CHEM 1A03/1AA3.

Originally Posted by RememberTwce

This site will be your friend.

Click "Orbitals" from the top menu and hover over an element. It has a bunch of stuff that will help for CHEM 1A03/1AA3.

You should've posted that last year.

thank you!!
now the second part of the question is to "Identify the general outer electron configuration for each group of elements shown in this periodic table"
how do I do that?

Originally Posted by Joyceyblank

Can someone check for me if this is correct? I got a few wrong on my first try, so here's the new one.

cesium- 6s^1
rutherfordium- 7s^2 6d^2
cadmium- 5s^2 4d^10
silicon- 3s^2 3p^2
ununhexium 7s^2 7d^10 7p^4
lawrencium 7s^2 5f^14 6d^10 7p^4
xenon 5s^2 4d^10 5o^6

thanks!

Rutherfordium you're missing the f electrons, lawrencium you're way off (only 1 d election and no p electrons) , otherwise it looks right from first glance. Are they really doing transition metal electron configurations in first year now?

Originally Posted by Joyceyblank

thank you!!
now the second part of the question is to "Identify the general outer electron configuration for each group of elements shown in this periodic table"
how do I do that?

Use the Aufbau "Building up" principle or in other words, fill the lowest energy orbitals first.

So electron shell 1 before 2 before 3 etc. (This is the n quantum number) Of course what you learn is that this is only for single electron species. When you get multiple electrons there is variation in energy between the subshells (s (0) vs. p (1) vs. d (2) vs. f (3) ) (This is the l quantum number). l can have a value of 0 to n-1 for a given shell. That means electron shell 1 has only s subshell, while 3 has s, p and d subshells. For multielectron atoms, d orbitals are higher in energy than the s orbital one shell lower and f orbitals are higher in energy than the f orbitals two shells lower.

So using the Aufbau principle fill 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p in that order. Look at your periodic table, now look back at me, now back at it. See how each row of the periodic table starts with two elements (ok 1 element in the first row, but that's cause shell 1 don't have p orbitals)? That's your s-block. After the first row the six columns at the far right are the p block. The awkward columns in the middle are the d-block. The black sheep that aren't in the actual table are the f-block. Ignore the d and f blocks for now. If you're filling an atom in the s or p block, assume all previous subshells are filled and just count the number of columns into the block it is from the left and that gives you the number of s or p orbitals respectively.


Another quantum number (ml) describes the orbital within the subshell. It can take integer values between -l and l. So there are 5 d orbitals (-2, -1, 0, 1, 2). Finally there is ms, which tells you the spin of the electron (+1/2 or -1/2). Two electrons can share an orbital but they are of opposite spins. So remember in a given subshell you can have twice as many electrons as there are orbitals, remember this when you're filling stuff out.

I apologize for the long sojourn into quantum numbers but they're useful for what i'm going to tell you next.

D and F blocks are weird. They don't follow the same rules precisely although its a good basis to start doing them like the s and p and then memorize exceptions. They assume strange arrangements to minimize energy. Half filled subshells (each orbital with one electron, aligned the same way) are particularly stable (and so are filled subshells). The stabilization occurring in the d and f subshells is enough to transfer an electron to or from the s or d subshells respectively. Also despite 4f being lower energy than 5d (and 5f lower than 6d) for the most part, its actually favourable for an electron to first join the d subshell. An f electron adds in the next element up and the next one up the d goes to an f. There are other exceptions too that don't follow the above explanations. tl;dr you'll have to memorize the d and f blocks but many of the strange configurations are explainable.

Hope that helps.

Originally Posted by RememberTwce

This site will be your friend.

Click "Orbitals" from the top menu and hover over an element. It has a bunch of stuff that will help for CHEM 1A03/1AA3.

This has a silly error. Zn, Cd, Hg are not transition metals. They are d block but are not classified as transition metals. But yes, I love this site, I forgot that I had it bookmarked.

Originally Posted by arathbon

Use the Aufbau "Building up" principle or in other words, fill the lowest energy orbitals first.

So electron shell 1 before 2 before 3 etc. (This is the n quantum number) Of course what you learn is that this is only for single electron species. When you get multiple electrons there is variation in energy between the subshells (s (0) vs. p (1) vs. d (2) vs. f (3) ) (This is the l quantum number). l can have a value of 0 to n-1 for a given shell. That means electron shell 1 has only s subshell, while 3 has s, p and d subshells. For multielectron atoms, d orbitals are higher in energy than the s orbital one shell lower and f orbitals are higher in energy than the f orbitals two shells lower.

So using the Aufbau principle fill 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p in that order. Look at your periodic table, now look back at me, now back at it. See how each row of the periodic table starts with two elements (ok 1 element in the first row, but that's cause shell 1 don't have p orbitals)? That's your s-block. After the first row the six columns at the far right are the p block. The awkward columns in the middle are the d-block. The black sheep that aren't in the actual table are the f-block. Ignore the d and f blocks for now. If you're filling an atom in the s or p block, assume all previous subshells are filled and just count the number of columns into the block it is from the left and that gives you the number of s or p orbitals respectively.


Another quantum number (ml) describes the orbital within the subshell. It can take integer values between -l and l. So there are 5 d orbitals (-2, -1, 0, 1, 2). Finally there is ms, which tells you the spin of the electron (+1/2 or -1/2). Two electrons can share an orbital but they are of opposite spins. So remember in a given subshell you can have twice as many electrons as there are orbitals, remember this when you're filling stuff out.

I apologize for the long sojourn into quantum numbers but they're useful for what i'm going to tell you next.

D and F blocks are weird. They don't follow the same rules precisely although its a good basis to start doing them like the s and p and then memorize exceptions. They assume strange arrangements to minimize energy. Half filled subshells (each orbital with one electron, aligned the same way) are particularly stable (and so are filled subshells). The stabilization occurring in the d and f subshells is enough to transfer an electron to or from the s or d subshells respectively. Also despite 4f being lower energy than 5d (and 5f lower than 6d) for the most part, its actually favourable for an electron to first join the d subshell. An f electron adds in the next element up and the next one up the d goes to an f. There are other exceptions too that don't follow the above explanations. tl;dr you'll have to memorize the d and f blocks but many of the strange configurations are explainable.

Hope that helps.

O_O Please tell me you copied and pasted that from somewhere.

Originally Posted by RememberTwce

O_O Please tell me you copied and pasted that from somewhere.

No, I figured it was good review for my Transition Metal Chem class.

what you said made sense, but why are the choices like this:
A. ns^2(n-1)d^5np^4
B. ns^1
C. ns^2(n-1)f^7
D. ns^2(n-2)f^14(n-1)d^10np^1
E. ns^2np^6
F. ns^2(n-2)f^14(n-1)d^11
G. ns^2(n-1)p^6
H. ns^2nd^1
I. ns^2(n-1)d^10np^4
J. ns^2(n-1)d^1
K. ns^2(n-1)d^7
L. ns^3
M. ns^2np^5
N. ns^ 2(n-2)f^

to match these highlighted blocks:
1. H, Li, Na, K, Rb, Cs, Fr
2. Sc, Y
3. Tl, Uut
4. Se, Te
5. Ne, Ar
6. Gd, Cm

thanks so much!

Originally Posted by Joyceyblank

what you said made sense, but why are the choices like this:
A. ns^2(n-1)d^5np^4

This wouldn't exist (it won't match anything). n-1 d subshell fills before the n p subshell. (n is the first quantum number, a d subshell of n-1 fills immediately after the ns subshell if it exists (d's only as part of shells 3 and up), and an f subshell of n-2 fills immediately after the ns subshell if it exists (f's only as part of shells 4 and up). Follow your periodic table. Move from left to right in a period and you should see s block then f block then d block then p block.

B. ns^1

One s electron would mean the first column (group) from the left of the s block. So this is the configuration of Hydrogen, Lithium, Potassium etc.

C. ns^2(n-1)f^7

You have a full s block, so move to the f block and move seven columns in. This is a half-filled f subshell so IIRC the elements in this column have what appears to be a normal electron configuration (not one of the exceptions I mentioned). Generally full or half full d (5 or 10 electrons) of f (7 or 14 electrons) follow the same rules as s and p blocks.

D. ns^2(n-2)f^14(n-1)d^10np^1

This is the first p block column. However since the f and d electrons are present you know its only the elements in one of the last two rows (where the f subshell fills)

E. ns^2np^6

This is a full s and p subshells. This would be the furthest right column in the p block, but only including the elements which aren't in a row with d or f electrons.

F. ns^2(n-2)f^14

Full s shell plus 14 columns into the f block.

G. ns^2(n-1)p^6

s and p subshells fill at the same level. So this is wrong.

H. ns^2nd^1

This doesn't happen. D subshell fills up after the s subshell one energy level up. (it should always say n-1)

I. ns^2(n-1)d^10np^4

Fourth column into the p block, the rows where they are proceeded by a d-block but not an f-block

J. ns^2(n-1)d^1

It should be the first d block element in rows where it isn't proceeded by an f-block.

K. ns^2(n-1)d^7

7th column into the d block, where they aren't proceeded by an f-block.

L. ns^3

Quantum comes in handy here. s is l=0. Therefore ml can only have one value (0) and thus there can only be one orbital or two electrons in any given s subshell.

M. ns^2np^5

Fifth column into the p block, where it isn't proceeded by the d or f blocks.

N. ns^ 2(n-2)f^

I think the last number was cut off, but its however many elements into the f-block.

Why has MI suddenly become the new Avenue? This is starting to get annoying...