Electronic configuration

Last updated: 31/01/2020

WJEC A-Level Chemistry Unit 1.6 – The periodic table Electronic configuration

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Question 1 of 100

1. Question
Why does the group containing the Nobel gases contain the highest ionisation energy compared to other groups?
- They have a full outer shell of electrons and high nuclear attraction
- They have an empty shell of electrons
- They have a full outer shell of electrons and low nuclear attraction
- They have an empty outer shell of electrons and no nuclear attraction
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Question 2 of 100

2. Question
What is the overall trend in ionisation energy across a period?
- Ionisation energy increases
- Ionisation energy decreases
- Ionisation energy slightly decreases
- Ionisation energy stays the same
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Question 3 of 100

3. Question
Explain the ionisation energy trend across a period.
- The number of protons in the nucleus increases, so there is more nuclear attraction. Electrons are added to the same shell, so the atomic radius decreases. There is the same number of shells so shielding will hardly change. Therefore, the ionisation energy will increase as you go along a period.
- The number of protons in the nucleus decreases, so there is less nuclear attraction. Electrons are added to the same shell, so the atomic radius decreases. There is the same number of shells so shielding will hardly change. Therefore, the ionisation energy will increase as you go along a period.
- The number of protons in the nucleus decreases, so there less nuclear attraction. Electrons are removed from the same shell, so the atomic radius increases. There is the same number of shells so shielding will hardly change. Therefore, the ionisation energy will decrease as you go along a period.
- The number of protons in the nucleus increases, so there is more nuclear attraction. Electrons are removed from the same shell, so the atomic radius increases. There is the same number of shells so shielding will hardly change. Therefore, the ionisation energy will increase as you go along a period.
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Question 4 of 100

4. Question
What happens to the ionisation energies between groups 2-13?
- Ionisation energy decreases slightly
- Ionisation energy decreases significantly
- Ionisation energy increases slightly
- Ionisation energy increases significantly
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Question 5 of 100

5. Question
Explain the pattern change in ionisation energies between groups 2-13.
- The outer electrons in group 13 are in p-orbitals. The outer electrons in group 2 are in s-orbitals. p-orbitals have a higher energy than s-orbitals so are further away from the nucleus and easier to remove so the ionisation energy decreases.
- The outer electrons in group 13 are in p-orbitals. The outer electrons in group 2 are in s-orbitals. s-orbitals have a higher energy than p-orbitals so are further away from the nucleus and easier to remove so the ionisation energy increases.
- The outer electrons in group 13 are in p-orbitals. The outer electrons in group 2 are in s-orbitals. s-orbitals have a higher energy than p-orbitals so are further away from the nucleus and easier to remove so the ionisation energy decreases.
- The outer electrons in group 13 are in p-orbitals. The outer electrons in group 2 are in s-orbitals. p-orbitals have a higher energy than s-orbitals so are further away from the nucleus and easier to remove so the ionisation energy increases.
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Question 6 of 100

6. Question
What happens to the ionisation energies between groups 1 and 2?
- Ionisation energy decreases sharply
- Ionisation energy increases sharply
- Ionisation energy slightly decreases
- Ionisation energy slightly increases
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Question 7 of 100

7. Question
Explain the pattern change in ionisation energies between groups 1 and 2.
- A new shell is added, increasing the distance of the outer shell from the nucleus and increasing the electron shielding.
- A new shell is removed, decreasing the distance of the outer shell from the nucleus and decreasing the electron shielding.
- A new shell is added, decreasing the distance of the outer shell from the nucleus and decreasing the electron shielding.
- A new shell is removed, increasing the distance of the outer shell from the nucleus and increasing the electron shielding.
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Question 8 of 100

8. Question
What is the overall trend in ionisation energy down a period?
- Ionisation energy decreases
- Ionisation energy increases
- Ionisation energy slightly decreases
- Ionisation energy stays the same
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Question 9 of 100

9. Question
Explain the ionisation energy trend down a group.
- The number of outer electron shells increases, so the distance of the electrons from the nucleus increases as well as the amount of shielding. Therefore, there is a weaker nuclear attraction. This outweighs the increasing nuclear charge, causing the ionisation energy to decrease.
- The number of outer electron shells decreases, so the distance of the electrons from the nucleus decreases as well as the amount of shielding. Therefore, there is a stronger nuclear attraction. This outweighs the increasing nuclear charge, causing the ionisation energy to decrease.
- The number of outer electron shells decreases, so the distance of the electrons from the nucleus increases as well as the amount of shielding. Therefore, there is a weaker nuclear attraction. This outweighs the increasing nuclear charge, causing the ionisation energy to increase.
- The number of outer electron shells increases, so the distance of the electrons from the nucleus decreases as well as the amount of shielding. Therefore, there is a weaker nuclear attraction. This outweighs the increasing nuclear charge, causing the ionisation energy to increase.
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Question 10 of 100

10. Question
Which of the following has the highest first ionisation energy? Be, Mg, Ca, Sr, Ba, Ra
- Be
- Ba
- Ra
- Ca
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Question 11 of 100

11. Question
Which of the following has the lowest first ionisation energy? Be, Mg, Ca, Sr, Ba, Ra
- Ra
- Ca
- Be
- Ba
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Question 12 of 100

12. Question
Which of the following has the highest first ionisation energy? B, C, N, O, F
- F
- O
- B
- N
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Question 13 of 100

13. Question
Which of the following has the lowest first ionisation energy? B, C, N, O, F
- B
- O
- F
- N
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Question 14 of 100

14. Question
Which of the following has the highest first ionisation energy? Rn, Xe, Kr, Ar
- Ar
- Rn
- Xe
- Kr
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Question 15 of 100

15. Question
Which of the following has the lowest first ionisation energy? Rn, Xe, Kr, Ar
- Rn
- Xe
- Ar
- Kr
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Question 16 of 100

16. Question
Which of the following has the highest first ionisation energy? K, Ca, Ga, Ge, As, Se, Br, Kr
- Kr
- K
- Ca
- Se
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Question 17 of 100

17. Question
Which of the following has the lowest first ionisation energy? K, Ca, Ga, Ge, As, Se, Br, Kr
- K
- Kr
- Ca
- Se
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Question 18 of 100

18. Question
Which of the following has the highest first ionisation energy? Rb, Fr, Li, K
- Li
- Fr
- K
- Rb
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Question 19 of 100

19. Question
Which of the following has the lowest first ionisation energy? Rb, Fr, Li, K
- Fr
- Li
- K
- Rb
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Question 20 of 100

20. Question
Which of the following has the highest first ionisation energy? Al, Si, P, S, Cl
- Si
- Cl
- S
- P
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Question 21 of 100

21. Question
Write the electronic configuration of calcium in its simplest form.
- [Ar] 4s $^2$
- [Ar] 4s $^1$
- [Ne] 4s $^2$
- [Ne] 4s $^1$
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Question 22 of 100

22. Question
Write the electronic configuration of bromine in its simplest form.
- [Ar] 4 $s^23d^{10}4p^5$
- [Ne] 1s $^22s^22p^63s^2p^64s^23d^{10}4p^5$
- [Ne] 4s $^23d^{10}4p^5$
- [Ar] 3s $^23d^{10}4p^5$
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Question 23 of 100

23. Question
Write the electronic configuration of magnesium in its simplest form.
- [Ne] 3s $^2$
- [Ar] 4s $^2$
- [Ne] 4s $^2$
- [Ar] 3s $^2$
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Question 24 of 100

24. Question
Write the electronic configuration of fluorine in its simplest form.
- [He] 2 $s^2p^5$
- [Ne] 2s $^2p^5$
- [He] 3s $^3p^5$
- [Ne] 3s $^3p^5$
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Question 25 of 100

25. Question
Write the electronic configuration of zinc in its simplest form.
- [Ar] 4s $^23d^{10}$
- [Ne] 4s $^23d^{10}$
- [Ar] 3s $^23d^{10}$
- [Ar] 4s $^24d^{10}$
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Question 26 of 100

26. Question
Write the electronic configuration of sulfur in its simplest form.
- [Ne] 3s $^23p^4$
- [Ar] 3s $^23p^4$
- [Ne] 4s $^23p^4$
- [He] 3s $^23p^4$
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Question 27 of 100

27. Question
Write the electronic configuration of oxygen in its simplest form.
- [He] 2s $^22p^4$
- [Ar] 2s $^22p^4$
- [He] 3s $^23p^4$
- [Ne] 2s $^22p^4$
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Question 28 of 100

28. Question
Write the electronic configuration of lithium in its simplest form.
- [He] 2s $^1$
- [He] 3s $^1$
- [Ne] 2s $^1$
- [Ne] 3s $^1$
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Question 29 of 100

29. Question
Write the electronic configuration of titanium in its simplest form.
- [Ar] 4s $^23d^2$
- [Ne] 4s $^23d^2$
- [Ar] 3s $^23d^2$
- [He] 4s $^23d^2$
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Question 30 of 100

30. Question
Write the electronic configuration of germanium in its simplest form.
- [Ar] 4s $^23d^{10}4p^2$
- [Ne] 4s $^23d^{10}4p^2$
- [Ar] 3s $^23d^{10}4p^2$
- [Ar] 4s $^23d^{10}3p^2$
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Question 31 of 100

31. Question
Write the electronic configuration of a bromine ion in its simplest form.
- [Ar] 4s $^23d^{10}3p^6$
- [Ar] 4s $^23d^{10}4p^5$
- [Ar] 3s $^23d^{10}4p^4$
- [Ar] 3s $^23d^{10}3p^6$
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Question 32 of 100

32. Question
Write the electronic configuration of a lithium ion in its simplest form.
- [He]
- [Ne]
- [Ne] 2s $^2$
- [He] 2s $^1$
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Question 33 of 100

33. Question
Write the electronic configuration of a potassium ion in its simplest form.
- [Ar]
- [Ne]
- [He] 3s $^1$
- [Ar] 4s $^1$
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Question 34 of 100

34. Question
Write the electronic configuration of an oxygen ion in its simplest form.
- [He] 2s $^22p^6$
- [Ne] 2s $^22p^2$
- [Ne] 2s $^22p^6$
- [He] 2s $^22^2$
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Question 35 of 100

35. Question
Write the electronic configuration of a chlorine ion in its simplest form.
- [Ne] 3s $^23p^6$
- [Ne] 3s $^23p^4$
- [He] 3s $^23p^6$
- [He] 3s $^23p^4$
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Question 36 of 100

36. Question
Write the electronic configuration of a magnesium ion in its simplest form.
- [Ne]
- [Ne] 3s $^2$
- [Ne] 3s $^22p^2$
- [Ne] 3s $^22p^3$
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Question 37 of 100

37. Question
Write the electronic configuration of a calcium ion in its simplest form.
- [Ar]
- [Ar] 3s $^2$
- [Ar] 3s $^23p^2$
- [Ar] 3s $^23p^3$
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Question 38 of 100

38. Question
Write the electronic configuration of a fluorine ion in its simplest form.
- [He] 2s $^22p^6$
- [Ar] 2s $^22p^6$
- [He] 2s $^23p^6$
- [Ar] 3s $^23p^6$
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Question 39 of 100

39. Question
Write the electronic configuration of a sulfur ion in its simplest form.
- [Ne] 3s $^23p^6$
- [He] 3s $^23p^2$
- [He] 3s $^23p^6$
- [Ne] 3s $^23p^2$
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Question 40 of 100

40. Question
Write the electronic configuration of an aluminium ion in its simplest form.
- [Ne]
- [Ne] 3s $^23p^6$
- [Ne] 3s $^23p^4$
- [He] 3s $^23p^4$
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Question 41 of 100

41. Question
What is ionisation?
- When atoms lose or gain electrons
- When atoms lose electrons
- When atoms gain electrons
- When atoms share electrons
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Question 42 of 100

42. Question
What is ionisation energy?
- Energy needed to form positive ions
- Energy needed to form negative ions
- Energy needed to form positive and negative ions
- Energy needed to create an ionic bond
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Question 43 of 100

43. Question
What is the definition of the first ionisation energy?
- The energy required to lose one mole of gaseous electrons from one mole of the gaseous element to form one mole of gaseous 1+ ions.
- The energy required to gain one mole of gaseous electrons to each atom in one mole of the gaseous element to form one mole of gaseous 1- ions.
- The energy required to gain one mole of electrons to each atom in one mole of the element to form one mole of 1- ions.
- The energy required to lose one mole of electrons to each atom from one mole of the element to form one mole of 1- ions.
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Question 44 of 100

44. Question
Why are outer shell electrons the first to be removed?
- Outer electrons experience the least nuclear attraction, as they are the furthest away from the nucleus and require the least ionisation energy.
- Outer electrons experience no nuclear attraction, requiring no ionisation energy.
- Outer electrons experience the most nuclear attraction, as they are the furthest away from the nucleus and require the most ionisation energy.
- The experience the most shielding.
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Question 45 of 100

45. Question
Explain how the atomic radius impacts the nuclear attraction experienced by an electron.
- The larger the atomic radius the smaller the nuclear attraction of the electrons.
- The smaller the atomic radius the smaller the nuclear attraction of the electrons.
- The larger the atomic radius the larger the nuclear attraction of the electrons.
- The atomic radius has no overall impact on the nuclear attraction experienced by an electron.
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Question 46 of 100

46. Question
Explain how the nuclear charge impacts the nuclear attraction experienced by outer electrons.
- The higher the nuclear charge, the larger the nuclear attraction on the outer electrons.
- The lower the nuclear charge, the larger the nuclear attraction on the outer electrons.
- The higher the nuclear charge the, the smaller the nuclear attraction on the outer electrons.
- The nuclear charge does not impact the nuclear attraction.
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Question 47 of 100

47. Question
What is shielding?
- The repulsion between inner and outer electron shells
- The attraction between inner and outer electron shells
- The shells that electrons occupy
- The repulsion from the nucleus
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Question 48 of 100

48. Question
Explain how the electron shielding impacts the nuclear attraction experience by outer electrons.
- The more inner shells there are, the larger the shielding effect is so the nuclear attraction experienced is smaller.
- The amount of shielding does not impact the nuclear attraction.
- The more inner shells there are, the lower the shielding effect is so the nuclear attraction experienced is smaller.
- The less inner shells there are, the larger the shielding effect is so the nuclear attraction experienced is larger.
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Question 49 of 100

49. Question
What are successive ionisation energies?
- A measure of the amount of energy required to remove each electron in turn.
- The energy required to remove one electron.
- The energy required to remove one mole of gaseous electrons from one mole of the gaseous element to form one more of gaseous 1+ ions.
- The energy required to gain one mole of gaseous electrons to each atom in one mole of the gaseous element to form one mole of gaseous 1- ions.
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Question 50 of 100

50. Question
What is second ionisation energy?
- A measure of how easily a +1 ion loses an electron to form a +2 ion.
- A measure of how easily a +2 ion gains an electron to form a +1 ion.
- A measure of how easily a +2 ion loses an electron to form a +3 ion.
- A measure of how easily a +1 ion gains an electron to form a stable ion.
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Question 51 of 100

51. Question
How many successive ionisation energies does lithium have?
- 3
- 1
- 2
- 4
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Question 52 of 100

52. Question
Write the first ionisation equation for lithium.
- $Li(g) -> Li^+(g) + e^-$
- $Li^+(g) -> Li^{2+}(g) + e^-$
- $Li^{2+}(g) -> Li^{3+}(g) + e^-$
- $Li^{-1}(g) -> Li(g) + e^-$
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Question 53 of 100

53. Question
Write the second ionisation equation for lithium.
- $Li^+(g) -> Li^{2+}(g) + e^-$
- $Li(g) -> Li^+(g) + e^-$
- $Li^{2+}(g) -> Li^{3+}(g) + e^-$
- $Li^{-1}(g) -> Li(g) + e^-$
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Question 54 of 100

54. Question
Write the third ionisation equation for lithium.
- $Li^{2+}(g) -> Li^{3+}(g) + e^-$
- $Li(g) -> Li^+(g) + e^-$
- $Li^+(g) -> Li^{2+}(g) + e^-$
- $Li^{-1}(g) -> Li(g) + e^-$
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Question 55 of 100

55. Question
Describe the trend in successive ionisation energies.
- Each ionisation energy is higher than the one before
- Each ionisation energy is lower than the one before
- The ionisation energy does not change
- Each ionisation energy is double the value of the one before
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Question 56 of 100

56. Question
What happens to ionisation energies when removing an electron from a different shell than beforehand?
- There is a large increase in energy
- They begin to decrease
- Nothing
- They drop to the first ionisation energy
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Question 57 of 100

57. Question
What lies on each axis when plotting an ionisation energy graph?
- Ionisation energy – ionisation number
- Ionisation energy – nuclear attraction
- Nuclear attraction – ionisation number
- Ionisation number – nuclear charge
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Question 58 of 100

58. Question
Write the first ionisation equation for sodium.
- $Na(g) -> Na^+(g) + e^-$
- $Na^+(g) -> Na^{2+}(g) + e^-$
- $Na^{2+}(g) -> Na^{3+}(g) + e^-$
- $Na^{-1}(g) -> Na(g) + e^-$
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Question 59 of 100

59. Question
Write the second ionisation equation for sodium.
- $Na^+(g) -> Na^{2+}(g) + e^-$
- $Na(g) -> Na^+(g) + e^-$
- $Na^{2+}(g) -> Na^{3+}(g) + e^-$
- $Na^{-1}(g) -> Na(g) + e^-$
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Question 60 of 100

60. Question
Write the third ionisation equation for sodium.
- $Na^{2+}(g) -> Na^{3+}(g) + e^-$
- $Na(g) -> Na^+(g) + e^-$
- $Na^+(g) -> Na^{2+}(g) + e^-$
- $Na^{-1}(g) -> Na(g) + e^-$
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Question 61 of 100

61. Question
Describe the movement of electrons.
- Orbiting the nucleus
- Stationary
- Random movements around an atom
- Moving within the nucleus
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Question 62 of 100

62. Question
What is a quantum number?
- Describes electrons in atoms
- Number of electrons
- Ratio between the number of protons and neutrons
- Number of protons
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Question 63 of 100

63. Question
What is the principle quantum number?
- Shows which shells are occupied by electrons
- Highest number of electrons
- Ion charge
- Number of shells
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Question 64 of 100

64. Question
How many electrons are in the first electron shell?
- 2
- 18
- 32
- 8
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Question 65 of 100

65. Question
How many electrons are in the second electron shell?
- 8
- 18
- 32
- 2
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Question 66 of 100

66. Question
How many electrons are in the third electron shell?
- 18
- 8
- 32
- 2
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Question 67 of 100

67. Question
How many electrons are in the fourth electron shell?
- 32
- 8
- 18
- 2
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Question 68 of 100

68. Question
What is the definition of an orbital?
- A region of space where electrons can be found
- A shell where electrons lie
- Shells that orbit the nucleus
- A region of space where protons can be found
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Question 69 of 100

69. Question
What shape is an s-orbital?
- Spherical
- Dumbbell
- Cuboid
- Oval
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Question 70 of 100

70. Question
What shape is a p-orbital?
- Dumbbell
- Spherical
- Cuboid
- Oval
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Question 71 of 100

71. Question
How many electrons can fill a s-orbital?
- 2
- 6
- 10
- 14
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Question 72 of 100

72. Question
How many electrons can fill a p-orbital?
- 2
- 6
- 10
- 14
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Question 73 of 100

73. Question
How many electrons can fill a d-orbital?
- 2
- 6
- 10
- 14
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Question 74 of 100

74. Question
How many electrons can fill a f-orbital?
- 2
- 6
- 10
- 14
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Question 75 of 100

75. Question
In terms of electronic structure, what does one box represent?
- One orbital
- One shell
- Number of shells
- An electron
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Question 76 of 100

76. Question
What property stops two electrons repelling when in an orbital?
- Spin
- Negative charge
- Nuclear attraction
- Orbiting around the nucleus
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Question 77 of 100

77. Question
In terms of electronic structure, how are electrons represented in the box formation?
- One arrow facing up, one arrow facing down
- Two arrows facing up
- Two arrows facing down
- One arrow facing up
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Question 78 of 100

78. Question
What is the charge of an electron?
- -1
- -2
- 2
- 1
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Question 79 of 100

79. Question
Explain the term shell.
- Energy level of the principle quantum number
- Energy level of quantum number
- Orbit the nucleus
- Holds electrons
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Question 80 of 100

80. Question
What is the maximum number of electrons an orbital can hold?
- 2
- 1
- 3
- 4
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Question 81 of 100

81. Question
What is an electron shell made up of?
- The same type of orbitals
- Orbiting electrons
- Empty space
- Electrons
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Question 82 of 100

82. Question
Describe the relationship between electrons occupying sub-shells and energy levels.
- Electrons occupy sub-shells in order of increasing energy levels
- Electrons occupy sub-shells in order of decreasing energy levels
- Electrons occupy sub-shells in random energy levels
- Electrons occupy sub-shells with no relationship to energy levels
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Question 83 of 100

83. Question
What is the definition of electron configuration?
- The arrangement of electrons in an atom or ion
- The arrangement of protons and neutrons
- The number of electrons
- The arrangement of electrons in a compound
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Question 84 of 100

84. Question
Which energy level is filled with electrons first?
- The energy level with the lowest available energy level
- The energy level with the highest available energy level
- The energy levels are filled in a random order
- The energy level furthest away from the nucleus
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Question 85 of 100

85. Question
Describe the relationship between energy levels in 4s and 3d orbitals.
- 4s orbital has a slightly lower energy level than the 3d orbital
- 4s orbital has the same energy level as a 3d orbital
- 4s orbital has a slightly higher energy level than the 3d orbital
- 4s orbital has a significantly higher energy level than the 3d orbital
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Question 86 of 100

86. Question
Write the electronic configuration of a bromine atom.
- 1s $^22s^22p^63s^23p^64s^23d^{10}4p^5$
- 1s $^22s^22p^63s^23p^64s^23d^{10}$
- 1s $^22s^22p^63s^23p^64s^23d^{10}4p^3$
- 1s $^22s^22p^63s^23p^64s^23d^{10}4p^4$
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Question 87 of 100

87. Question
Write the electronic configuration of a magnesium atom.
- 1s $^22s^22p^63s^2$
- 1s $^22s^23s^2$
- 1s $^22s^22p^6$
- 1s $^22s^22p^63s^1$
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Question 88 of 100

88. Question
Write the electronic configuration of a carbon atom.
- 1s $^22s^22p^2$
- 1s $^22s^2$
- 1s $^22p^2$
- 1s $^22s^22p^3$
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Question 89 of 100

89. Question
Write the electronic configuration of a fluorine atom.
- 1s $^22s^22p^5$
- 1s $^22s^22p^4$
- 1s $^22p^5$
- 1s $^22s^22p^6$
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Question 90 of 100

90. Question
Write the electronic configuration of a nickel atom.
- 1s $^22s^22p^63s^23p^64s^23d^8$
- 1s $^22s^22p^63s^23p^64s^24d^8$
- 1s $^22s^22p^63s^23p^63d^8$
- 1s $^22s^22p^63s^23p^64s^23d^7$
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Question 91 of 100

91. Question
Write the electronic configuration of a titanium atom.
- 1s $^22s^22p^63s^23p^64s^23d^2$
- 1s $^22s^22p^63s^23p^63d^5$
- 1s $^22s^22p^63s^23p^64s^24d^2$
- 1s $^22s^22p^63s^23p^64s^2$
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Question 92 of 100

92. Question
Write the electronic configuration of a lithium atom.
- 1s $^22s^1$
- 1s $^22s^2$
- 1s $^22s^22p^63s^1$
- 1s $^22s^22p^63s^2$
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Question 93 of 100

93. Question
Write the electronic configuration of a sodium atom.
- 1s $^22s^22p^63s^1$
- 1s $^22s^22p^63s^2$
- 1s $^22s^2$
- 1s $^22s^22p^6$
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Question 94 of 100

94. Question
Write the electronic configuration of an oxygen atom.
- 1s $^22s^22p^4$
- 1s $^22s^22p^63s^4$
- 1s $^22s^22p^5$
- 1s $^2$
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Question 95 of 100

95. Question
Write the electronic configuration of a lithium ion.
- 1s $^2$
- 1s $^1$
- 1s $^22s^1$
- 1s $^22s^2$
Correct

Incorrect
Question 96 of 100

96. Question
Write the electronic configuration of a chlorine ion.
- 1s $^22s^22p^63s^23p^6$
- 1s $^22s^22p^63s^23p^5$
- 1s $^22s^22p^63s^23p^7$
- 1s $^22s^22p^6$
Correct

Incorrect
Question 97 of 100

97. Question
Write the electronic configuration of a calcium ion.
- 1s $^22s^22p^63s^23p^6$
- 1s $^22s^22p^63s^23p^64s^2$
- 1s $^22s^22p^63s^23p^64s^23d^2$
- 1s $^22s^22p^6$
Correct

Incorrect
Question 98 of 100

98. Question
Write the electronic configuration of an oxygen ion.
- 1s $^22s^22p^6$
- 1s $^22s^22p^63s^23p^6$
- 1s $^22s^22p^63s^23p^4$
- 1s $^22s^4$
Correct

Incorrect
Question 99 of 100

99. Question
Write the electronic configuration of a potassium ion.
- 1s $^22s^22p^63s^23p^6$
- 1s $^22s^22p^63s^23p^64s^2$
- 1s $^22s^22p^63s^23p^64s^2$
- 1s $^22s^22p^6$
Correct

Incorrect
Question 100 of 100

100. Question
Write the electronic configuration of a sodium ion.
- 1s $^22s^22p^6$
- 1s $^22s^22p^63s^23p^6$
- 1s $^22s^22p^63s^2$
- 1s $^22s^22p^63s^1$
Correct

Incorrect