Part C: Concept Checks

Transcript

0:00Introduction

Hello everyone! Okay, one quick statement before we start. Whenever an atom gains an electron, energy is released. Is that statement true or false? Chlorine releases energy when it gains an electron, and so does oxygen. So the statement must be true, right? Hold that thought. We will settle it in Example two. This is Part C of the unit problem bank: four concept checks on periodicity, two quick ones first and then two worked examples, the kind of ideas that are often tested in statement form, where a single word can change the answer. So, let's begin.

0:38Two enthalpies, two signs

First, the two quantities behind today's checks. Ionisation enthalpy, delta i H: the Greek capital delta, a small i for ionisation, and H for enthalpy. Take an isolated atom, X, as a gas; the small g in brackets means gas. It loses one electron, the e with a minus sign, and the arrow gives the gaseous positive ion, X plus. Okay? Then electron gain enthalpy, delta e g H, with a small e g below the delta for electron gain. Now the atom gains one electron instead, and we get the gaseous negative ion, X minus. Right? And the sign convention for both of them. Negative means energy is released. Positive means energy is taken in, so it has to be supplied. Keep these two signs in mind, because both our examples depend on them. Clear?

1:32Check 1: do isotopes have different radii?

Now for check one, on isotopes. Isotopes of an element have the same atomic number, capital Z, the number of protons. But they have different masses, because they carry different numbers of neutrons. Okay? Now let me ask you this. The heavier isotope has a heavier nucleus. So is the heavier isotope also a bigger atom? No, it is not, and here is why. The radius of an atom is set by its electrons, and by the pull they feel. That pull is Z effective, the capital Z with a small e f f below it, which equals Z minus S, where S is the screening by the other electrons. Extra neutrons add mass, but no charge. With me? Isotopes have the same Z, so they have the same nuclear charge, and the same electrons in the same arrangement. Right? So isotopes of an element have the same radius, for practical purposes. Any difference is far too small to matter. Got it? And here is how it comes in statement form. Chlorine thirty five minus is smaller than chlorine thirty seven minus; that is the less-than sign. The small number at the top left is the mass number, and the raised minus is the charge of the chloride ion. This statement is false, because isotopes have equal size. Clear?

2:52Check 2: where does hydrogen belong?

Check two, and look at this outline of the periodic table. The periods, one to seven, run down the side, and the group numbers run across the top. The s, p, d and f blocks are shaded. Now see the single box at the top left, in period one, group one. Okay? That box is for hydrogen, H, with the configuration one s one, which means a single electron in the one s subshell; the small raised one counts it. So by its configuration, with its one s electron, hydrogen goes in group one, in that box. Right? And it behaves like group one, the alkali metals. Like them, it has one outer electron, and it can lose it to form H plus, the positive ion. With me? Now count the other way. The first shell holds only two electrons. So how many electrons is hydrogen short of a full shell? Yes, exactly, just one more electron completes its shell! And being one electron short of a full shell is the mark of group seventeen, the halogens, in the column just before group eighteen. Got it? So hydrogen also behaves like a halogen. It gains an electron to form H minus, the hydride ion, and it forms H two, a molecule of two atoms; the small two below counts the atoms. Okay? So hydrogen fits both families. It has one electron, like the alkali metals, and it is one electron short, like the halogens. That is exactly why it is hard to place in the periodic table. Clear?

4:25Example 1: ionisation enthalpy is never negative, even Na

Now, two examples from this part of the problem bank. Example one, an easy one to warm up. Is ionisation enthalpy ever negative? Pause the video and try it first. Done? Okay, let's check. Sodium gives up its outer electron very easily. So for sodium, maybe removing the electron releases energy, and the value is negative, right? No, no, we have to be careful here, so ask what holds the electron. The nucleus attracts the electron it holds, in every atom, and that includes sodium. Okay? So to pull the electron away, we must work against that attraction. Removing an electron always needs energy, so delta i H is greater than zero; that is the greater-than sign on screen. Right? Look at sodium, N a, in group one: it has the lowest first ionisation enthalpy in its period. Yet delta i H one, with the small one for the first electron, is still plus four hundred and ninety six kilojoules per mole; the minus one on mol means per mole. Easy to remove means small, not negative. Got it? So the answer is no. Ionisation enthalpy is never negative; it is always positive. With me?

5:44Example 2: positive electron gain enthalpy, noble gases, O2-

And now our statement from the start, as Example two, which has more cases to remember. Can electron gain enthalpy be positive? Pause the video and try it first. Here is a hint: think about atoms that are already very stable. Done? Okay, let's check. For most elements, delta e g H is less than zero; the less-than sign means it is negative, so energy is released. That is the part our opening statement got right. Okay? But most is not all. So tell me, which atoms would not welcome an extra electron? Yes, it is the noble gases, very good! Helium, plus forty eight; neon, plus one hundred and sixteen; and argon, plus ninety six, all in kilojoules per mole. These values are positive, so energy has to be supplied. Right? The reason is simple: their subshells are already full. So the added electron must enter the next shell, far from the nucleus and well shielded. With me? Two more cases give positive or near zero values. Group two elements have a full n s two subshell, where n is the outer shell, and the small raised two counts its two electrons. And nitrogen, N, has a half-filled two p subshell, which is extra stable. Got it? And one more case, for every element. Suppose an atom has already gained one electron, and become a negative ion. Does that negative ion welcome a second electron? Exactly, it does not, because the negative ion repels the incoming electron! Take oxygen: O minus, plus an electron, gives O two minus; the raised two minus is its charge. This second step needs plus eight hundred and forty four kilojoules per mole. And every second electron gain is like this: energy is taken in. Clear? So the answer is yes. Electron gain enthalpy can be positive for the noble gases, for group two elements and nitrogen, and for every second electron gain. So our opening statement, that energy is always released, is false! Got it?

7:59Quick revision: both enthalpy signs, isotopes and hydrogen

So, quick revision. First, ionisation enthalpy is always positive, because removing an electron always needs energy. Second, electron gain enthalpy can be positive: for the noble gases, group two, nitrogen, and every second electron gain. Third, isotopes have the same radius, and hydrogen is hard to place, because it is like group one and also like group seventeen. Now go and try the practice problem on this page. Read every word of the question. All the best!