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Chemistry library
Course: Chemistry library > Unit 8
Lesson 2: Periodic table trends- Periodic trends
- Atomic radius trends on periodic table
- Atomic and ionic radii
- Mini-video on ion size
- Ionization energy trends
- Ionization energy: period trend
- First and second ionization energy
- Electron affinity: period trend
- Electronegativity
- Electronegativity and bonding
- Metallic nature
- Periodic trends and Coulomb's law
- Worked example: Identifying an element from successive ionization energies
- Ionization energy: group trend
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Worked example: Identifying an element from successive ionization energies
When electrons are removed in succession from an element, the transition from removing valence electrons to removing core electrons results in a large jump in ionization energy. By looking for this large jump in energy, we can determine how many valence electrons an element has, which in turn can help us identify the element. Created by Sal Khan.
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I get the idea of needing more energy to remove the core electrons. But once you remove the electron from the 2nd shell, wouldn't the next electron be easier to remove. For that matter, shouldn't the first 3 removals take less and less energy, and then have a big spike to remove the 4th electron, and then less to remove the fifth?(5 votes)
No. Each successive ionisation energy is always greater than the previous one.(2 votes)
I haven't received and answer yet for my question 3 days ago, and this is the revised version:
Why can't the identity of the element be P? Though s and p orbitals are in the same shell, the p orbital is supposed to have higher energy. Please check my reasoning below:
1. When you add 3p1, the outermost electron is in the 3p orbital (even though 3s and 3p are all valence electrons).
2. Therefore the 3p1 electron is subject to electron-electron repulsion by 3s2 electrons
3. The distance between the nucleus and the electron in 3p is further than the 3s2 electrons
By using 1,2,3 in Coulombs's law, when you add 3p1, ionization energy decreases. Which also makes sense because a lone electron farthest away from the nucleus should be easiest to pull away.
In reverse therefore, when you take away 3p1, ionization energy increases, and that could account for the leap between the third and fourth IE.
Thank you in advance for your answer.(4 votes)
When we look at the changes between different ionization energies, we see an abnormally large increase from the third to the fourth ionization energies. Much higher than the transition from any of the other ionization energies. Since ionization energy is the amount of energy required to remove an electron, a high ionization indicates a difficult to remove electron. This tells us the atom resisted losing that fourth electron more so than the first three. The reason for this resistance is because the atom reaches a relatively stable electron configuration by losing the first three electrons and does not want to disturb that stability by losing a fourth electron.
Such a large difference between the third and fourth ionization energy corresponds to the stability associated with a noble gas electron configuration. So we’re looking for a third period element which has a cation of +3 charge with the same electron configuration as the previous noble gas, neon. This can only be aluminum then.
There is an energy difference between the 3s and 3p subshells, but it’s slight and this wouldn’t correspond to such a large energy difference as the one between the third and fourth ionization energies here. The main culprit for huge energy difference between ionization energies is the cation achieving a noble gas electron configuration and not wanting to lose that stability.
For phosphorus we would expect a large difference between the fifth and sixth ionization energies, not between the third and fourth. To confirm this here are the first seven ionization energies of phosphorus in kJ/mol: 1011.8, 1907, 2914.1, 4963.6, 6273.9, 21267, 25431; and we can see that abnormally large difference between the fifth and sixth ionization energies as expected. And like aluminum, its +5 cation has the same electron configuration as neon and wants to remain as such.
Side note, you don’t need to post the same question multiple times to get a response. People will respond when they have time, not because you repeat it. Doing so makes you look needy.
Hope that helps.(2 votes)
- If the period of the element isn't given, how would we determine what element it is?(2 votes)
- If you narrow down the options to a group 13 element then you would need to know the ionization energies of that group’s elements, or their relative amounts. Ionization energy within a group generally decreases as you move from top to bottom. This is because as you move down the valence electrons are located in shells farther away from the nucleus and so feel less attraction to the nucleus and are easier to remove. This means elements like boron and aluminum will have similar ionization energy patterns which show them to be in the same group, but boron will have a higher first ionization energy than aluminum.
So without actually providing the ionization energies for all the group 13 elements, they could say that the element has the second highest first ionization energy in its group, which would be aluminum. But even that wouldn’t work well since gallium (the element beneath aluminum) has about the same first ionization energy as aluminum. The general rule is that ionization energy decreases down a group, but the pattern isn’t as identifiable as the pattern across a period.
Most questions wouldn’t bother with all this nuance and would simply just say which period the element is in.
Hope that helps.(4 votes)
- What about a third period element with the numbers:
IE1 1230 kJ/mol
IE2 2400 kJ/mol
IE3 3600 kJ/mol
IE4 8000 kJ/mol
IE5 10500 kJ/mol
IE6 22000 kJ/mol
IE7 27500 kJ/mol
IE8 35200 kJ/mol(1 vote)
The biggest jump (change) is between IE5 and IE6 so I would guess that from IE6 & up we are removing core electrons.
This means that we most likely have 5 valence electrons, meaning we would be in the 15th group on the periodic table, meaning we would guess Phosphorous (P).
I'm not sure if you were challenging me or asking a question but I hope this helps & is beneficial to someone. :-)(4 votes)
- Why can't the identity of the element be P? Though s and p orbitals are in the same shell, the p orbital is supposed to have higher energy. Please check my reasoning below:
1. When you add 3p1, the most outer electron is in the 3p orbital (even though 3s and 3p are all valence electrons).
2. Therefore the 3p1 electron can be shielded by 3s2 electrons
3. The distance between the nucleus and the electron in 3p is further than the 3s2 electrons
By 1,2,3, when you add 3p1, ionization energy decreases. Which also makes sense because a lone electron farthest away from the nucleus should be easiest to pull away.
In reverse therefore, when you take away 3p1, ionization energy increases, and that could account for the leap between the third and fourth IE.
Thank you in advance for your answer.(3 votes)- Although the 3p orbital has a higher energy level than the 3s orbital, the 3s orbital does not shield the 3p. The 3p orbital has higher energy because of the warped nature of the p-orbital, not because the 3p orbital is any further from the nucleus than the 3s orbitals. Additionally, adding 3p1 would not decrease the ionization energy. When you add valence electrons, ionization energy generally increases because of a decreased shielding effect.(0 votes)
- Why can't Boron or Gallium be an acceptable answer as well?(2 votes)
- "The first five ionization energies for a third-period element are shown below", key word being third-period element. Boron and Gallium have the same trend, but are in the second and fourth periods respectively. Additionally the numbers given are only true for aluminium. Boron and Gallium have different values.
Hope that helps.(2 votes)
Why is there a difference between the first, second, and third ionization energies? For the first three electrons removed, they were all valence electrons so shouldn't they all have been equally as easy to remove?(1 vote)- Subsequent ionization energies of all elements have a general increasing trend because it becomes progressively harder to remove additional electrons. When we remove the first electron from a neutral atom with the first ionization energy, it becomes a cation, or a positively charged ion. Having an overall positive charge means the remaining electrons feel a greater force of attraction to the atom. Essentially by removing one electron, the remaining electrons have one less repulsive force (because electrons repel each other with their negative charge) but the same attractive force (from the positive protons) resulting in a net increase of attractive force. Removing even more electrons will result in even higher third, fourth, fifth, etc. ionization energies because the magnitude of the positive charge continues to increase.
Hope that helps.(3 votes)
- Can we say N is atom that have a same ionazitaion energy, because it show same ionazitaon ratio, or we vant say because N can not make 5.th ionazition?(1 vote)
- Can you rephrase that? It doesn't make sense.(3 votes)
- So I understand that it has three valence electrons, but wouldn't the other elements in that column also have 3? so why is Al?(1 vote)
- The question indicated that the element must be a third-period element. Therefore the element must be in the third row. Hope this helps!(1 vote)
Video transcript
- [Instructor] We are
told that the first five ionization energies for
a third period element are shown below. What is the identity of the element? So pause this video and see if you can
figure it out on your own and it'll probably be handy to have a periodic table of elements. So before I even look at a
periodic table of elements, let's make sure we
understand what this table is telling us. This is telling us that if
we start with a neutral atom of this mystery element, it would take 578 kilojoules per mole to remove that first electron to turn that atom into an ion with a plus one positive charge. And then, it would take another
1,817 kilojoules per mole to remove a second electron. So to make that ion even more positive. And then after that it would take another 2,745 kilajoules per mole to
remove the third electron. And then to remove the fourth electron, it takes a way larger amount of energy. It takes 11,578 kilojoules per mole. And then the fifth
electron takes even more, 14,842 kilojoules per mole. And so, for the first, second, and third you do have an increase
in ionization energy, but when you go to the
fourth the energy required to remove those is way higher. So to me, these look like you're
removing valence electrons and these look like you're
removing core electrons. So one way to think about it is let's look on our
periodic table of elements and look for a third period element that has three valence electrons. So we have our periodic table of elements. We want a third period element, so it's gonna be in this third row and which of these has
three valence electrons? Well, sodium has one valence electron, magnesium has two valence electrons, aluminum has three valence electrons. So one way to think about it is that first electron, it's a
reasonable ionization energy. Then the second one, a little higher. Then the third, a little
bit higher than after that, but then the fourth, you're
starting to go into the core. You're going to have to take an electron out of that full second energy shell, which takes a lot of energy. And so this is pretty clearly aluminum that is being described.