The Electron Configuration: Ions - Online Tutor, Practice Problems & Exam Prep

When elements lose or gain electrons, they form cations and anions, respectively. Cations are positive ions, and when they lose electrons, it occurs from orbitals with the highest principal quantum number (n-value), which indicates the shell number or energy level. For example, electrons are removed from 3p before 3s because it is higher in energy. Anions, on the other hand, are negatively charged ions that gain electrons. These electrons are added to orbitals with available space. In naming anions, the nonmetal base name is retained but ends with "-ide" (e.g., oxide, fluoride, chloride). Understanding the process of electron removal or addition and the resulting ion forms is crucial for grasping ionic compounds and their electron configurations.

The Electron Configuration: Ions

Recall that when elements lose or gain electrons, they become ions. We're going to say here that cations are positive ions, and we're going to say what that cat ion. We first remove electrons from the highest shell number, and the highest shell number is connected to our principal quantum number, or the N value.

Now the principal quantum number provides the shell number, or energy level of the electron. So here we have 1S22S22P63S23P6. The number is related to the principal quantum number. N1S2 would have an N value of 12S2 and 2P6 would have N values equal to two. 3S2 and 3P6 would have N values equal to three.

So just remember we look at the number before the sub level or subshell letters of SPD or F to determine our N value, and when we're removing electrons, we're going to remove them from the orbitals that have the highest N value. But let's say we wanted to remove electrons and we only had 3S2 and 3P6 to choose from. Both of them have the highest end value.

But we'd say that since 3P6 is listed after 3S2, it would have more energy and therefore if we had to choose between 3S2 and 3P6, we'd remove the electron first from 3P6. Same thing can be said between 2S2 and 2P6. Since 2P6 is listed after 2S2, it has more energy and therefore if we had to take electrons away, it'd come out of the 2P6 orbitals before we even look at the 2S2 orbitals.

So just remember, first look at the highest end value. If a set of orbitals have the same end value, then the higher energy one is the one that was listed last. That's where we would remove electrons first.

The Electron Configuration: Ions Example 1

So here it says write the condensed electron configuration for the titanium 3 ion. Alright, so titanium three is TI-3. Plus, step one tells us to provide the electron configuration for the neutral form of the element. All right, so titanium neutral has an atomic number of 22. If we're doing its condensed electron configuration. Looking at the periodic table, we would see argon as a noble gas right before it. Then we'd have here 4 S 2, and then we'd have 3-D2. So this would be the electron configuration of the neutral titanium atom.

Step 2 Begin removing electron or electrons from the highest numbered shell to obtain the desired charge. When in the same shawl 2S versus 2P use off bowel principle to remove the higher energy electrons first. All right, so coming back to the titanium, 3 ion 3 plus means we've lost 3 electrons. Looking at our condensed electron configuration for the neutral ion, a neutral element will help guide us to the new electron configuration of the ion.

So we don't touch what's within these brackets. That's our noble gas. So we're looking here because it's 4S, that means N = 4, and because it's 3D, that means N = 3. We remove our electrons first from the highest shell number, so we need to lose 3 electrons here. We're going to lose our first two from the 4S, so 4S is now completely gone. Next we need to lose one more electron because we need to lose 3 electrons, so that's going to come from the 3D. That means titanium three at the end will be Argon 3D ONE as the electron configuration of our ion.

The Electron Configuration: Ions

Now for an anion. Remember anions are negatively charged ions. We become anions when we accept electrons. We're going to say with an anion you add an electron or electrons to the orbitals with available space.

Now note for an anion, the nonmetal keeps its base name but has its ending changed to eyed. So if we hear oxide, fluoride, chloride, that means we're dealing with the anion form of that element. So keep that in mind when asked to find the electron configurations of anions.

The Electron Configuration: Ions Example 2

Here it says write the full electron configuration for the nitride ion, so it ends with I'd, which means we're dealing with the anion form of nitrogen. So nitrogens in Group 3A, so its charge is -3 or 3 -. 3 minus means that we have gained 3 electrons, so we'll take that into account when we're doing its electron configuration.

So the steps we're going to do here is we're going to provide the electron configuration for the neutral form of the element. So nitrogen when it's neutral, it has an atomic number of seven. Its full electron configuration is 1S22S22P3. Next Step 2 we're going to add electron or electrons to the orbitals that can accommodate more electrons. So for the nitride ion, here we have this S can only hold up to two electrons Max, SO1S2 and 2S2 stay the way that they are.

Remember, the P sublevel can hold up to six electrons. Right now it only has three. We're gaining three additional electrons, so all of them can be tossed into this 2P6. So this would represent the full electron configuration of the nitride ion. So we can say this as an acceptable answer.

One more thing we can also say is that it now has the same electron configuration as a noble gas, and a noble gas would be neon. So both of these answers are acceptable. You can show the full electron configuration here and you can say that it is the same Electron configuration as Neon. Technically in this answer I'm asking for the full electron configuration, so this would be the best answer, but just keep in mind it now also has the same electron configuration as a noble gas, so if we wanted to do a shorthand or condensed electron configuration you could just write this as an answer.

What is the full electron configuration of the selenide ion?

• A. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴
• B. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶
• C. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴ 5s²
• D. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁸ 4p⁴ 5s² 5p²

Determine the electron configuration for the Cl⁺ ion.

• A. [Ne] 3s¹ 3p⁵
• B. [Ne] 3s² 3p⁶
• C. [Ne] 3s² 3p⁵
• D. [Ne] 3s² 3p⁴

Determine the electron configuration and electron orbital diagram for the silver ion.

• A. [Kr] 4d¹⁰
• B. [Kr] 4d⁸ 5p²
• C. [Kr] 4d⁸
• D. [Kr] 4d⁸ 5s²

Here’s what students ask on this topic:

How do you write electron configuration for ions?

Writing electron configurations for ions involves the same principles as for neutral atoms, but you must account for the gain or loss of electrons due to the ion's charge. Here's how to do it:

1. Determine the number of electrons in the ion. For cations (positively charged ions), subtract the charge from the atomic number to find the number of electrons. For anions (negatively charged ions), add the charge to the atomic number.
2. Write the electron configuration for the neutral atom. Use the periodic table to determine the order in which the orbitals fill up (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.).
3. Adjust the electron configuration for the ion. If it's a cation, remove electrons starting from the highest energy level (outermost shell). If it's an anion, add electrons to the next available orbital in the sequence.

For example, for Na⁺ (sodium ion), the neutral sodium atom has 11 electrons, but Na⁺ has lost one electron, leaving it with 10 electrons. The electron configuration for Na⁺ is the same as that for neon (Ne), a noble gas: 1s² 2s² 2p⁶. For Cl^- (chloride ion),

What is the term for two different ions with the same electron configuration?

The term for two different ions with the same electron configuration is "isoelectronic." Isoelectronic species are atoms, molecules, or ions that have the same number of electrons and hence the same electronic structure. This can occur despite them having different numbers of protons or being different elements altogether. A classic example of isoelectronic ions would be the sodium cation (Na⁺) and the neon atom (Ne), both of which have ten electrons and thus share the same electron configuration of:

${\mathrm{1s}}^2{\mathrm{2s}}^2{\mathrm{2p}}^6$

However, they differ in their nuclear charges, with sodium having 11 protons and neon having 10 protons. Isoelectronic species often exhibit similar chemical and physical properties due to their analogous electron arrangements.