EARTH |
Yeah, yeah. Pauling's electronegativity Table shows why. Oxygen's electronegativity is 3.44 and hydrogen's is 2.20, so the difference (1.24) means that the electrons will be shared between them most of the time. But a few of them will be ionic bonds. We don't know which are ionic and which are covalent, just that most are covalent. |
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A good guess and the right one. But remember that you MUST calculate the difference in electronegativity in order to figure out the bonds. Don't go by instinct (unless you must!).OK. OK. The difference in electronegativity for HF is 1.78. That's probably a big enough difference to make an ionic bond. Right? |
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OK, Lithium (Li) has an electronegativity of 0.98 while hydrogen's
electronegativity is 2.20. A difference of 1.22. Hey, that's not
enough to form an ionic bond. That's less than 1.5!
Now Arthur! What did I say about that number? |
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You have the wisdom of Solomon, and you are not alone. Most Alchemists display hydrogen somewhere in between. I like to move it closer to the Group VII elements because that places it near the elements that form covalent bonds too. I just remind myself that hydrogen is nothing more than a proton! Usually.But hydrogen may also have some neutrons. |
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Yeah, metals form metallic bonds. Their outer shell electrons
are kind of loose and they form a "sea of electrons"
instead of specific bonds. This "super sharing" of outer
shell electrons is a metallic bond.
Yes, very good. The most important metal property is the ability to conduct electrons.It's that "super sharing" of electrons. That allows metals to pass electrons along very quickly, producing a current of electrons.Right, but there are other metal properties. Remember malleability and ductility?Oh yeah, I forgot.Metals are malleable, meaning they can be hammered or pressed into thin sheets. And they are ductile, meaning they can be drawn into a wire (by pulling it through a hole). |
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OK. (The hydrogen will just remind me we're on Group I.)
All Group I elements have a single electron in their outer
shell and form cations readily.
I've never thought of sodium as a metal. I just think of it as
a salt.
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You said they do this quickly. Are all the Group I elements equally
quick.
No. The speed and efficiency of this reaction changes. Their reactivity increases as you go down Group I. When exposed to air, lithium (Li) takes on a slight tarnish and does so rather slowly. Sodium is quicker. This reactivity continues to increase (speed up) as you go down Group I. Pure cesium (Cs) will react so quickly that when placed in moist air it actually fizzes.Like sodium when placed in water! What happens if you drop a chunk of cesium into a bowl of water?What do you think would happen?I bet it would react VERY quickly.You are right. Dropping cesium into water causes a massive
explosion!
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It means that they also form excess hydroxide ions (OH-) and have good metallic behavior. But their reactivity and metallic behaviors are not as good as Group I.So Group II elements have the same important properties as Group I, but do not produce as much hydroxide as Group I or conduct electricity as well as Group I. Right?Yes. Group II elements are kind of "mini" alkali metals. Instead of calling them "mini-alkali-metals", we call the Group II elements "alkaline metals". The "ne" at the end reminds us they are less than a full alkali metal (Group I).Kind of confusing isn't it. I mean the Group I are more alkaline, but we call them "alkali". However, the Group II elements are less alkaline but we call them simply "alkaline". Why have Alchemists made it so backward? |
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That's right! All the Group II elements react less violently than their neighbor on the left. Beryllium (Be) is less reactive than lithium (Li). Magnesium (Mg) is less reactive than sodium (Na). This is true of all the comparisons of the first and second Groups.So they studied the Group II elements before the Group I elements because it was easier!Aye. That's it. The Group II elements are like their neighbors to their left (Group I) but not so extreme.I see. Do the same trends continue down Group II? Like reactivity and softness increasing as you go down the Group II elements?Yes, they do. Beryllium (Be) is the hardest and least reactive of the Group II elements. Magnesium (Mg) is softer and more reactive than beryllium, but magnesium is harder and less reactive than the element below it, calcium (Ca). The trend continues. |
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Let me try this.
I would guess that calcium (Ca) is more reactive than magnesium because calcium is below magnesium. But calcium is less reactive than potassium (K) or strontium (Sr) because those elements are to the left (K) or below (Sr) the calcium. Right? Right! By Jove I think you've got it!It's kind of like a staircase. As you go down or to the left the reactivity increases.That is an excellent observation, Arthur. You will see staircases a lot in the Table. As a matter of fact, we see staircases as we continue to move from right to left across the Table. |
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Look at this Periodic Table. I've drawn it to show the metals, nonmetals and semimetals.The staircase if obvious. It looks like a staircase of metal with a carpet of semimetals and then a layer of all the other elementsYes, that is exactly the way it looks to me too. You will notice that elements in Group III and beyond are either metals, nonmetals or semimetals.You know, this staircase seems to start just after the Group II elements. It's as if something special happens there. Isn't this where those missing elements belong? You know, the gap in atomic number between calcium (atomic number 20) and gallium (atomic number 31). |
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The only element among the post-transitional metals that I recognize
is lead (Pb as in "plumber"). That's the metal used
to make pipes! Right?
Right. In our ancient world, lead was the only post-transitional metal to be mined. It is easy to work with and easy to make into pipes.So it's the most important post-transitional metal!Well, yes, in our medieval times. However, aluminum (Al) is the most important metal in the 20th and 21st centuries. Aluminum is the most abundant metal on earth. Or on the moon! It is the best metal in the world for manufacturing and construction. |
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That's right. Let's get back to the Table and continue our
tour.
I suppose that is the effect of this staircase structure. They
are not fully metals or nonmetals. They are in between.
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I'm not surprised.
So the elements called semimetals in the Table have the semi-electrical
properties we talked about, but not necessarily the chemical properties?
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Oh yeah. Carbon's a semimetal.
Hey! You just said it was a nonmetal! Are you making this up as you go? No, I'm not! Carbon is an amazing element. It sits in a very important part of the Table, where all the important properties of elements seem to come together. I keep coming back to the "fuzziness" of this part of the Table because it gets complex.Sure does!Carbon, like all the Group IV elements, has four electrons in its outer shell. Both carbon and silicon can share these outer electrons to form covalent bonds. |
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With only five electrons in their outer shell they need three
more to complete it. They would probably make three covalent bonds.
Ah, but....EXCEPT (you didn't let me finish!),Except for the two metals further down (antimony, Sb and bismuth, Bi) which would form metal bonds by "super sharing". And then we must not forget that awful semimetal (arsenic, As) which makes me dizzy just thinking about it! Yes, you are right. Sorry about your dizziness, Arthur.
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Are you going to say "ah well, sort of" and then tell
me something to confuse me?
I'm really trying hard not to, Arthur.Let's have it. I can take it.OK. Just a minor point really. The four elements at the top of Groups V and VI have some peculiar tricks about their bonds. |
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Kind of like nitrogen.
There seems to be a "flip-flop" among these four elements at the top of Groups V and VI. How do you mean?Nitrogen is very inert, but the element right below it, phosphorous, is very reactive. At least as white sulfur. But in Group VI the reactive element is oxygen, at the top, and sulfur is the nonreactive element. Reactivity "flip-flops" among these four elements. |
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Because they form salts. "Halogen" is from the Greek meaning "salt-producer". The most common salt on Earth is sodium chloride (NaCl). Geologists call rocks of NaCl, "halite".The halogens sound important.They are. Except astatine (At). It's too rare. But all the others are fairly common, and they can all form covalently bonded, diatomic, elemental molecules.Which means they make F2, Cl2, Br2 and I2 by sharing electrons in their outer shell.Right. Like O2 except they make only one covalent bond between them........ because they need to share only one electron to complete their outer shell. |
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Wait! Wait! I'm lost.
I'm not surprised. The point to understand is that the Group
VII elements can form covalent elemental molecules. But
these elemental molecules are so reactive that (given the chance)
they will quickly react with what ever they can in order to become
an ionic compound.
Maybe.
When Group VII elements are all alone, nothing but fluorine for
example, then the only thing they can do is form covalent bonds
by sharing their outer electron.
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Think about what the electronegativity means to elements in those two Groups. The decreasing electronegativity of Group I means they are MORE likely to give up an electron to become a cation. The decreasing electronegativity of the Group VII means they are LESS likely to become anions.I see.The effects are opposite because they are creating opposite ions. |
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In the middle of the Table the atom can put electrons into the d and f orbitals of its second outermost shell. They kind of back up and use d and f orbitals that they skipped. |
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Yes, it is. Notice that once all the d orbitals of that inner shell are full the next element will behave correctly by placing its next electron in a p orbital of its outer shell. So that element is a typical element. |
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Like being hard and having a high melting point.
Right.Hey, what about mercury (Hg, atomic number 80)?!It's a liquid! It isn't as strong as iron. It isn't strong at all! |
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Probably.
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Yes, very good observation. What do you think that means?That you've hidden another Table, and my guess is that it contains even weirder elements! |
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Lanthanum (La) is the transition element of the sixth Period that "introduces" the inner transition elements of that Period. |
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OK. Here is the complete Periodic Table of the elements.Wow. The Table is very long. |
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Look at the Table again. It looks like big gaps and changes occur when the new orbitals are added. |
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Yes, I do. Congratulations, Arthur. You have discovered a wee bit about the Table I was not going to cover. You are discovering how the orbitals affect properties. That "hiccup" in the ionization energies as you went from the Group II to Group III elements is due to the addition of a new type of orbital. The p orbital. Can you guess why?Hmmmm...It should be harder to ionize a Group III element than a Group II element, because the electron in the outer shell should be attracted closer to the nucleus of a Group III element and held tighter to it. But it looks like that isn't true on this ionization Table. At least for the first three pairs (Be-B, Mg-Al, and Ca-Ga). It's as if the first electron to be put into a p orbital is held loosely. |
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Notice that all the fourth Period elements have an N-shell as their outer shell, but the M-shell is used by the transition elements of the fourth Period as a place to store their electrons (in d orbitals). |
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The inner transition elements start in the sixth Period. The lanthanoid series starts with cerium (Ce), the first element to use f orbitals. It ends with lutetium (Lu). The actinoid series is in the seventh Period and starts with thorium (Th). |
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We've covered a lot of information here. I think these last Tables prove it! Read through it all again and check your notes. Take a good look at the Tables. I'm going to draw my own Periodic Table.