configuration writing practice
write a ground state electron configuration for each neutral atom. ground state means that all of the lowest possible energy levels (up to the proper number of electrons for the element) are filled.
1. na 2. pb 3. sr 4. u 5. n 6. ag 7. tl 8. ce 9. cl 10. hg
write a ground state electron configuration for these ions. remember that ions have a change in the total number of electrons (positive have lost electrons and negative have gained).
example: n3- is 1s² 2s² 2p⁶. it has three extra electrons.
11. o²- 12. fe²+ 13. b³+ 14. ni²+ 15. k+ 16. co³+
an excited atom has an electron or electrons which are not in the lowest energy state. excited atoms are unstable energetically. the electrons eventually fall to a lower level. * is used to indicate an excited atom. for example: *li 1s² 3p¹. (the ground state for li is 1s² 2s¹.)
write an excited state electron configuration for each. 17. al 18. ar 19. k 20. c
21. if each orbital can hold a maximum of two electrons, how many electrons can each of the following hold? a. 2s b. 5p c. 4f d. 3d e. 4d
22. what is the shape of an s orbital?
23. how many s orbitals can there be in an energy level?
24. how many electrons can occupy an s orbital?
25. what is the shape of a p orbital?
26. how many p orbitals can there be in an energy level?
27. which is the lowest energy level that can have a s orbital?
28. which is the lowest energy level that can have a p orbital?
29. is it possible for two electrons in the same atom to have exactly the same set of quantum numbers?
30. distinguish between an atom in its ground state and an excited atom.
31. how many d orbitals can there be in an energy level?
32. how many d electrons can there be in an energy level?
33. which is the lowest energy level having d orbitals?
34. how many f electrons can there be in an energy level?
35. which is the lowest energy level having f orbitals?

Question

Accepted Answer

# Explanation:
## Step1: Recall electron - configuration rules
Use the Aufbau principle, Pauli - exclusion principle, and Hund's rule to write electron configurations.
## Step2: Determine atomic numbers
For neutral atoms, the number of electrons is equal to the atomic number. For ions, adjust the number of electrons according to the charge.
### 1. Na (Atomic number = 11)
$1s^{2}2s^{2}2p^{6}3s^{1}$
### 2. Pb (Atomic number = 82)
$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{2}$
### 3. Sr (Atomic number = 38)
$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}$
### 4. U (Atomic number = 92)
$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{6}7s^{2}5f^{3}6d^{1}$
### 5. N (Atomic number = 7)
$1s^{2}2s^{2}2p^{3}$
### 6. Ag (Atomic number = 47)
$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{1}4d^{10}$
### 7. Tl (Atomic number = 81)
$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{1}$
### 8. Ce (Atomic number = 58)
$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{1}5d^{1}$
### 9. Cl (Atomic number = 17)
$1s^{2}2s^{2}2p^{6}3s^{2}3p^{5}$
### 10. Hg (Atomic number = 80)
$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}$
### 11. $O^{2 -}$ (O has atomic number 8, $O^{2 -}$ has 10 electrons)
$1s^{2}2s^{2}2p^{6}$
### 12. $Fe^{2 +}$ (Fe has atomic number 26, $Fe^{2 +}$ has 24 electrons)
$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{6}$
### 13. $B^{3+}$ (B has atomic number 5, $B^{3+}$ has 2 electrons)
$1s^{2}$
### 14. $Ni^{2+}$ (Ni has atomic number 28, $Ni^{2+}$ has 26 electrons)
$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{8}$
### 15. $K^{+}$ (K has atomic number 19, $K^{+}$ has 18 electrons)
$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}$
### 16. $Co^{3+}$ (Co has atomic number 27, $Co^{3+}$ has 24 electrons)
$1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{6}$
### 17. Al (Atomic number = 13), excited state example: $1s^{2}2s^{2}2p^{6}3s^{1}3p^{2}$ (electron promoted from 3s to 3p)
### 18. Ar (Atomic number = 18), excited state example: $1s^{2}2s^{2}2p^{6}3s^{2}3p^{5}4s^{1}$ (electron promoted from 3p to 4s)
### 19. K (Atomic number = 19), excited state example: $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4p^{1}$ (electron promoted from 4s to 4p)
### 20. C (Atomic number = 6), excited state example: $1s^{2}2s^{1}2p^{3}$ (electron promoted from 2s to 2p)
### 21.
a. 2s: Since each orbital can hold 2 electrons, 2s can hold 2 electrons.
b. 5p: The p - subshell has 3 orbitals, so it can hold $3	imes2 = 6$ electrons.
c. 4f: The f - subshell has 7 orbitals, so it can hold $7	imes2=14$ electrons.
d. 3d: The d - subshell has 5 orbitals, so it can hold $5	imes2 = 10$ electrons.
e. 4d: The d - subshell has 5 orbitals, so it can hold $5	imes2=10$ electrons.
### 22. The shape of an s orbital is spherical.
### 23. There is 1 s orbital in an energy level.
### 24. An s orbital can hold 2 electrons.
### 25. The shape of a p orbital is dumb - bell shaped.
### 26. There are 3 p orbitals in an energy level.
### 27. The lowest energy level that can have a s orbital is n = 1.
### 28. The lowest energy level that can have a p orbital is n = 2.
### 29. No, according to the Pauli - exclusion principle, two electrons in the same atom cannot have exactly the same set of quantum numbers.
### 30. An atom in its ground state has all electrons in the lowest possible energy levels. An excited atom has one or more electrons in higher - energy levels than the ground - state configuration.
### 31. There are 5 d orbitals in an energy level.
### 32. The d - subshell has 5 orbitals, and each orbital can hold 2 electrons, so there can be 10 d electrons in an energy level.
### 33. The lowest energy level having d orbitals is n = 3.
### 34. The f - subshell has 7 orbitals, and each orbital can hold 2 electrons, so there can be 14 f electrons in an energy level.
### 35. The lowest energy level having f orbitals is n = 4.
# Answer:
1. $1s^{2}2s^{2}2p^{6}3s^{1}$
2. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{2}$
3. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}$
4. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}7s^{2}5f^{3}6d^{1}$
5. $1s^{2}2s^{2}2p^{3}$
6. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{1}4d^{10}$
7. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}6p^{1}$
8. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{1}5d^{1}$
9. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{5}$
10. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{6}6s^{2}4f^{14}5d^{10}$
11. $1s^{2}2s^{2}2p^{6}$
12. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{6}$
13. $1s^{2}$
14. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{8}$
15. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}$
16. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{6}$
17. $1s^{2}2s^{2}2p^{6}3s^{1}3p^{2}$ (example)
18. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{5}4s^{1}$ (example)
19. $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4p^{1}$ (example)
20. $1s^{2}2s^{1}2p^{3}$ (example)
21. a. 2; b. 6; c. 14; d. 10; e. 10
22. Spherical
23. 1
24. 2
25. Dumb - bell shaped
26. 3
27. n = 1
28. n = 2
29. No
30. Ground state: electrons in lowest energy levels; Excited state: electrons in higher energy levels
31. 5
32. 10
33. n = 3
34. 14
35. n = 4

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Question

Explanation:

Step1: Recall electron - configuration rules

Step2: Determine atomic numbers

1. Na (Atomic number = 11)

2. Pb (Atomic number = 82)

3. Sr (Atomic number = 38)

4. U (Atomic number = 92)

5. N (Atomic number = 7)

6. Ag (Atomic number = 47)

7. Tl (Atomic number = 81)

8. Ce (Atomic number = 58)

9. Cl (Atomic number = 17)

10. Hg (Atomic number = 80)

11. $O^{2 -}$ (O has atomic number 8, $O^{2 -}$ has 10 electrons)

12. $Fe^{2 +}$ (Fe has atomic number 26, $Fe^{2 +}$ has 24 electrons)

13. $B^{3+}$ (B has atomic number 5, $B^{3+}$ has 2 electrons)

14. $Ni^{2+}$ (Ni has atomic number 28, $Ni^{2+}$ has 26 electrons)

15. $K^{+}$ (K has atomic number 19, $K^{+}$ has 18 electrons)

16. $Co^{3+}$ (Co has atomic number 27, $Co^{3+}$ has 24 electrons)

17. Al (Atomic number = 13), excited state example: $1s^{2}2s^{2}2p^{6}3s^{1}3p^{2}$ (electron promoted from 3s to 3p)

18. Ar (Atomic number = 18), excited state example: $1s^{2}2s^{2}2p^{6}3s^{2}3p^{5}4s^{1}$ (electron promoted from 3p to 4s)

19. K (Atomic number = 19), excited state example: $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4p^{1}$ (electron promoted from 4s to 4p)

20. C (Atomic number = 6), excited state example: $1s^{2}2s^{1}2p^{3}$ (electron promoted from 2s to 2p)

21.

22. The shape of an s orbital is spherical.

23. There is 1 s orbital in an energy level.

24. An s orbital can hold 2 electrons.

25. The shape of a p orbital is dumb - bell shaped.

26. There are 3 p orbitals in an energy level.

27. The lowest energy level that can have a s orbital is n = 1.

28. The lowest energy level that can have a p orbital is n = 2.

29. No, according to the Pauli - exclusion principle, two electrons in the same atom cannot have exactly the same set of quantum numbers.

30. An atom in its ground state has all electrons in the lowest possible energy levels. An excited atom has one or more electrons in higher - energy levels than the ground - state configuration.

31. There are 5 d orbitals in an energy level.…

Answer: