What was Newlands describing in his "Law of Octaves"? What is the "rule" that Newlands found and what is its modern equivalent that we use in bonding? What is the importance of the number "eight"? |
Mendeleev developed the first Periodic Table of the elements and it described the behavior of most of the smaller elements but not all of the elements. Moseley found a different way to arrange the elements. What had Mendeleev done wrong and what was Moseley's discovery that set the Table right? |
What are the first 20 elements in order of increasing atomic number?
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In the Periodic Table, what is a Group and what is a Period? What do elements in the same Group have in common? What do elements in the same Period have in common? |
What is a "metalloid"? |
What name is used to describe the three forms of pure carbon (graphite, diamonds and fullerenes)? |
Name an element with two gaseous allotropes. |
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Alchemists use X-rays to measure the sizes of atoms. (You knew
that.)
Why are the atoms in table salt (NaCl) of different sizes to those
listed as atomic radii?
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Assume you had an element able to make covalent bonds and ionic bonds of either ion, as well as being all alone.
List, from smallest
to largest, how the size of that atom would change for each type
of atom and bond.
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What atom (element) has all the properties in question 13? That is, what element can form covalent bonds, ionic bonds of either ion and also be atomic? |
Which Table would you use to calculate the energy needed to make Na+ from Na? |
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Describe "electron affinity". |
The noble elements (Group VIII) have high ionization energies
and low electron affinities. What does that tell you about their
ability to form ions?
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What is "electronegativity" and how is it used by Alchemists? |
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The difference in electronegativity between hydrogen and fluorine is 1.78. Would you expect that HF to be held together by covalent bonds or electrovalent bonds? |
You may have been surprised by that last answer. (But I bet you get this one!) Is the H-F bond a polar bond? Can HF make hydrogen bonds? |
Hydrogen fluoride (HF) is a "hydrogen halide", meaning a compound of hydrogen and a halogen (Group VII). Name the other hydrogen halides and briefly explain how they compare to each other in terms of their polarized bonds. |
When you add hydrogen halides to water (by bubbling the gas through water) they break up into ions, forming "hydrohalic acids". The extra hydrogen ions added to the water cause the water to be described as an acid. (Acids have excessive amounts of hydrogen ions.) An Alchemist would write that as
Which hydrogen halide do you think would be the last one to release its hydrogen ion to the water? |
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Freon is usually a gas except when the molecules are pushed very close to each other. That is, freon becomes a liquid only if you push the molecules together with a lot of pressure. As soon as the pressure is released, the freon molecules fly away from each other and return to their natural state, a gas. Why? Why is freon so reluctant to form the intermolecular bonds with other freon atoms which would cause it to stay a liquid? What causes molecules of freon to push away from each other? |
Table salt (NaCl) dissolves in water but not in hexane. Why? |
Fingernail polish and most paints are made of non-polar compounds (with lots of C-H bonds). They don't wash off in water (easily), but they will come off if scrubbed with hexane. Why? |
Explain how sodium hydride (NaH) can be formed.
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Does the electronegativity of the Group I elements increase, decrease or remain the same when you go down the Group? And why? |
Does the reactivity of the Group I elements with water increase, decrease or remain the same when you go down the Group? And why? |
How do the electronegativity and reactivity with water of the Group I and Group II elements compare? And why? |
Describe the properties of the alkaline metals that gives them their name? |
Are the alkaline metals more alkaline than the alkali metals? |
Do the metallic properties of the Group I and II metals decrease down the Groups? And why (not)? |
Does the electronegativity of the halides (Group VII) increase or decrease as you go down the Group? |
Does the reactivity of the halides, as elemental molecules, increase or decrease down the Group? And why? |
If electronegativity decreases down all Groups,
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Do semimetals make semimetal bonds? Explain. |
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Here's another look at the complete Periodic Table of the Elements. How many electrons should the M, N, O, and P-shells be able to hold? |
8 protons, 8 neutrons, with its electrons arranged as 2 (in the K-shell) and 6 (in the L-shell). The L-shell electrons are the outer electrons and they will be arranged in orbitals as s2 and p4. (We'll ignore exactly which types of p orbitals the 4 electrons can be placed in.)
Use this method and this Table to describe 12C.
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Do the same (as in Q 49) for 40Ar.
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Calcium (Ca) is element number 20 and the last "typical" element before the transition elements start to make things complex. Scandium (Sc) is the first transition element (atomic number 21). Tests done on scandium prove that it has 2 electrons in its outer (N) shell, just like calcium!
Describe what the electron arrangement of these two atoms (Ca and Sc) should
be. Then try to figure out how scandium could have 2 electrons in its outer shell. Where does the extra electron go? |
What are the number of protons and neutrons in 137Ba? Also, how are barium's electrons distributed among its shells?
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Do the same (as in Q 53) for 197Au. You'll need to see the Tables clearly, so I've shown you two of them.
Take your time with this puzzle. Gold is an inner transition metal, so its electron structure is complicated, but try to predict its orbital configuration anyway. (It's good practice.) |
Let's take a break from discussing electrons (phew) and think a wee bit about neutrons.
Here's a reminder of the atoms and isotopes we've discussed recently.
Do you see a progression? Describe how the ratio of protons to neutrons changes as the atoms get larger. |
Alchemists sometimes include the atomic number and atomic mass with the atom's symbol in order to give the maximum amount of information. Alchemists use a "Z" to stand for atomic number (the number of protons) and place it as a subscript before the symbol. They also use an "A" for the atomic mass (the number of nucleons) and include it as a superscript before the symbol (like you learned earlier). If the element has the symbol "X" then the entire atom is written as AZX. Rewrite the atoms from question 54 in this manner so you can see their atomic mass (A) as a superscript and atomic number (Z) as a subscript to each symbol.
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OK, back to electron configuration (ugh).
As you now know it is very difficult to keep track of the electronic configuration of an atom. Some Alchemists have to figure our the electronic configuration of all the inner shells as well as the outer shells! They don't always have a book nearby to help them. Sometimes they don't even have a Table to guide them!
First, let's recall that the shells, which we've been calling "K-shell", "L-shell", etc. are really based upon the first quantum number. What is the (first) quantum number for each shell?
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Take a look at this chart. It is a simple list of the various orbitals for each shell. Notice I am now using the principle quantum number for the shells instead of their letters.
Draw this chart into your notebook. The first column lists all the s orbitals with their principal quantum number in front of each. The second column is a list of all the p orbitals also with the principal quantum number preceding it. (Notice that we start the second column with principal quantum number 2 because the second shell, the L-shell, is the first to have p orbitals.) The third column lists the d orbitals starting with principal quantum number 3 (because d orbitals start with the M-shell). The chart ends with the fourth column where you list the f orbitals starting with principal quantum number 4.
Draw the chart carefully, following the logic I explained above.
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(Once you have this all written out you will be surprised how easy this is.) |
Now that you have your orbital series worked out (that's the name for the sequence of orbitals all arranged in increasing order) you can easily work out the orbital configuration of any atom without looking at the Periodic Table or worrying about the complexities of "basement" shells. Use your newly gained ability to work out the complete electron orbitals to all these atoms you did earlier.
(Notice I've included their proton counts, Z number, as a superscript so you don't need to consult a Table.) |
Let's find the electron configuration of another three elements using the orbital energy series.
Try
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Alchemists use a shorthand to write electron orbitals. They use noble elements to keep track of things. That is, they compare the electron configurations to that of the noble element that preceded it. (Note, that means we are talking about the Group VIII element that is a Period BEFORE the atoms we are concerned with.) Find the electronic configuration of the noble element that precedes nickel, copper and zinc and use it as a foundation on which to "build" the nickel, copper and zinc electronic configurations. (Use copper's correct configuration.) |
What are the "transuranic elements"? Where are they hidden in the Table? And what is special (unusual) about them? |
List the four kinds of orbitals,
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Newlands noticed that when (some) of the elements are listed in ever increasing mass, they tended to repeat their properties every eighth time. He called this observation his "Law of Octaves". This is the evidence behind Lewis's "octet rule". We use Lewis's octet rule to draw Lewis structures in order to understand bonding. Atoms can hold eight electrons in their outermost shell. It is this desire to have a complete outer shell (that is an outer shell with eight electrons) that gives rise to most bonding and chemical behavior. Indeed, there can never be more than 8 electrons in the outermost shell. (Transition and inner transition elements hide their extra electrons in lower shells.) |
Mendeleev developed a theory about the chemical periodicity, but he arranged the elements according to their atomic mass, instead of their atomic number. This is pretty good for the smaller elements. Moseley's X-ray work allowed him to count the protons and arrange the elements into increasing atomic number. When Alchemists arranged the elements according to Moseley's atomic NUMBER (instead of atomic mass), the elements with similar electronic configuration lined up into patterns whose behavior repeated every eighth time. |
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In the Periodic Table each column (running from top to bottom) is called a Group and each row (running across) is a Period. Elements in the same Group have the same type of outer shell. Note that they do not necessarily have the same number of electrons in the outer shell (although they often do). Elements in the same Group have the same needs to complete their outer shell. Therefore elements within a Group often have similar chemical properties. By knowing the behavior of one element in a Group, you can make a good guess that the others will behave the same. Periods run from left to right in rows, each row repeating when the outer shell is full. Elements in the same Period have the same outer shell BUT not the same shell type. For example, the first row, or "Period", represents the K-shell and both elements in the first Period (hydrogen and helium) have a K-shell as their outer shell. (That's assuming they are not ionized.). The second period is the L-shell and all elements in the second Period have an L-shell as their outermost shell. |
A metalloid is a semimetal displaying a chemical property of a metal, like the ability to turn water alkaline. |
Graphite, diamonds and fullerenes are allotropes of carbon.
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Oxygen has two allotropes, "normal" oxygen (O2) which we need to breathe and ozone (O3) which we don't want to breathe, but which protects us from the harmful ultraviolet light. |
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This is another difficult problem and you may have been a bit frustrated with it. But I bet you learned something!
You have every reason to be confused. Out of frustration you may
have decided that radon must be 1.34 angstroms. But that is not
the correct way to look at it.
"So, what is the size of radon?", you ask? (Perhaps shouting!). I don't know. No one really knows for sure. Even the sizes of xenon and astatine are argued about. The reason for this disagreement is that the lower right corner of the Periodic Table includes elements that are very difficult to purify and measure. Many of these elements are rare and unstable (radioactive). Some books give sizes for these atoms but another book will give different sizes! That is because of differences in the methods they use to prepare and measure the elements. If you found this problem irritating, I am sorry. But I think you will have learned that some things are not as straight forward as we all would like them to be. However, the rules about how the size of atoms change as you move around the Table still apply. |
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The largest type of atom would be an anion. The extra electron may require a new shell to be formed, thus making the atom larger than "normal". If there is room in the outer shell for the extra electron, that shell will grow a wee bit larger because there are fewer protons than electrons, so the electrostatic attraction can't hold all the electrons as close. Summary: from smallest to largest; cation, covalent, "normal", and then anion. |
Hydrogen! (Of course!).
Note: the hydride (H-) will have the electronic shell of helium (a full K-shell), but will be larger than helium. That is because the hydride has only one proton pulling in two electrons and it doesn't do that very well. So the hydride has a bigger (K) shell than helium even though they both have two electrons in it.
Don't get confused. As you move across a Period (from left to
right) the atoms ALWAYS get smaller.
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You must use the Table of first ionization energies to figure out how much energy (electron volts, eV) is needed to ionize Na to Na+. |
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The second ionization energy decreases down the Group, just like the first ionization energy. The reason is the same as for the first ionization. Larger atoms have larger shells with more electrons. The larger shells mean the electrons in the outer shell are easier to remove than they would be if they were in the Period above it. The same trend is seen in second ionization energy as seen in the first ionization energy.
Read through that again to be sure you got it. |
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That was a VERY difficult question and you should be very proud of yourself if you got either answer. However, you should read through this answer and make sure you know what I have done to get the two possible answers. Oh, which is the right answer? Well, that is hard to say. A great deal depends on how the energy is delivered to each atom. Even an advanced Alchemist would have trouble proving which way these three atoms would end up! |
Electron affinity is a measure of how easily an atom accepts an extra electron to become an anion. It is NOT the opposite of ionization energy. By that I mean that an atom with low or even negative electron affinity is not necessarily good at becoming a cation. |
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Electronegativity is a measure of how tightly an atom holds onto
electrons - its own electrons or spare electrons it finds.
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Fluorine is the most electronegative of all elements and that makes it the most reactive of all elemental molecules. Because
fluorine has the greatest electronegativity of any element (3.98,
just to remind you) it can pull electrons away from just about
any molecule it approaches. The only atom safe around a fluorine
atom is another fluorine atom. Two fluorine atoms (F2) are covalently
linked together because they both have the exact same electronegativity,
so they can't steal an electron away from the partner. All they
can do is share a pair of electrons as they try to satisfy their
"desire" for a complete outer shell (8 electrons). But
when an atom of any other element comes along, fluorine tries
to grab an electron away from it. Most of the time the other atom
doesn't stand a chance! If the other element has a low electronegativity,
fluorine will steal an electron. That causes fluorine to become
an anion and the other atom to become a cation.
So to summarize, fluorine gas (F2) is only stable with itself, because only then can the tug-of-war between the two powerful atoms cancel each other out. But when that gas meets any other kind of atom it either steals an electron from it (if the "victim" has a very low electronegativity of its own), or forms new covalent bonds to it (if the "victim" has a reasonable electronegativity of its own). |
The hydrogen gas (H2) is a covalently linked molecule. The two hydrogen atoms share a pair of electrons because they have the exact same electronegativity (2.20). But when mixed with fluorine gas (which you described in the previous question) the electrons in the two hydrogen atoms cannot resist the pull of the fluorine. This causes the hydrogen molecule (H2) to break up as the fluorine steals away the electrons. The difference in electronegativity
between fluorine and hydrogen (3.98 - 2.20 = 1.78) proves that
hydrogen cannot put up a fight with those strong fluorines.
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The huge difference in electronegativity (1.78) is more than the
1.50 that I like to use as my "cut off". But that only
goes to illustrate that there is no clear cut off between covalent
and ionic bonds.
Recall the O-H bonds in water are mostly covalent but have some ionic bonding some of the time. And the O-H bonds have a difference in electronegativity of only 1.24. Water's slight ionic character is helped by oxygen's lone pair orbitals and other details involving its structure. HF has certain properties that affect its character too. H-F bonds, like O-H bonds, are covalent for the most part but have some electrovalent bonding some of the time. Surprise! As you work through the next questions you will see how this affects the covalent bond |
You will also have noticed that a very polarized bond like H-F will allow the hydrogen to form hydrogen bonds. HF's extremely polarized bond causes the hydrogen to be extremely delta plus, so it would easily form a hydrogen bond with any atoms that is even slightly negative. |
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Hydrogen fluoride has the highest electronegativity difference (1.78) among the hydrogen halides. So, it will be the last molecule to give up its hydrogen.
Before leaving the subject of hydrohalides and the acids they
form, it is worth repeating that HF does NOT make as strong an
acid as the other hydrogen halides (HCl for example). This confuses
some people.
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Recall that carbon has a covalency of four, meaning it forms 4 covalent bonds (when it can). Each halide has a covalency of one. So, freon is a carbon surrounded by the four halide atoms. (It's shaped like methane but with chlorines and fluorines instead of hydrogen.) Carbon is the least electronegative of the atoms in the freon molecule, so it does not hold its shared electrons as tightly as the chlorine and fluorine atoms. The carbon atom at the center of freon will have a partial positive charge (delta plus) and the halides will all have a partial negative charge (delta minus). |
Freon has four atoms sticking out from it which have slightly negative charges. These negative, partial charges are caused by the delta minus on the four halide atoms. The over all effect is to create a "shield" of partial negative charges around each freon molecule. This causes freon molecules to repel each other! Remember, molecules of the same charge will fly away from each other because of electrostatic repulsion (the exact opposite of electrostatic attraction). Therefore, you have to push very hard to make two freon molecules get close enough together to form any kind of intermolecular bonds (bonds between molecules). Actually, it is hard to imagine ANY intermolecular forces could form with all those partial negative charges on their surface. But it CAN be done if you put enough pressure on them. As soon as the pressure is released the molecules quickly fly away from each other, returning to their gaseous state. This property, along with some others that are beyond the scope of this class, made freon the molecule of choice in refrigerators. Freon is a very stable molecule, most of the time, and many folks came to think of it as inert. But ultraviolet light can break up freon and create ions of chlorine and fluorine that will attack and destroy ozone (O3). It is for this reason that nations decided late in the 20th century to stop releasing freon into the atmosphere. (There are lots of other things to be said about freon's chemistry and physics, but they are too advanced a subject for this course.) |
NaCl is an ionic compound, held together by electrovalent bonds.
These bonds are formed because there is an electrostatic attraction
between the sodium cation (Na+) and the chlorine anion (Cl-).
Polar and ionic molecules dissolve best in a polar liquid (water) but they do not dissolve in a non-polar liquid (hexane). |
Fingernail polish and most paints are non-polar, so the polar nature of water is no help at all. As a matter of fact, water's polar nature causes the water molecules to stick to each other and hardly touch any of the non-polar molecules. That is, water doesn't "interact" with non-polar molecules very well (hardly at all!).
But when you wipe hexane on fingernail polish it sweeps away the
layers of polish because...
Note and Warning! Hexane is a carcinogen (causes cancer)
and is flammable (burns easily). It should NOT be used on skin, nails or any other part of your body.
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(Remember the "cut-off" of 1.5 is not a true cut-off.) The NaH molecule is formed because sodium is not able to hold
onto its outer electron when hydrogen pulls on it. In this tug-of-war
the sodium gives up its outer electron to become a cation. (Nothing
unusual about that.) The hydrogen gains that electron and becomes
an anion (H-)! Then the two ions are drawn together by electrostatic
attraction to form the ionic molecule NaH.
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The electronegativity of ALL the elements (not just Group I) decreases down the Group, because the extra shells added with each Period make it easier to remove an electron and harder to hold onto an extra electron. |
The reactivity of the Group I elements with water increases down
the Group because with decreasing electronegativity, the water
molecules can more easily grab an electron away.
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The alkaline metals are alkaline metals!
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No, it's the other way around.
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No, it's the other way around.
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Electronegativity of all the elements decreases down the Groups (because the extra shells make the atoms larger). |
The reactivity of the halides decreases down the Group because the electronegativity also decreases as you go down (any) Group. |
The reactivity of Groups I and VII in water change in opposite directions because the way they react is opposite. Group I elements react in water by donating an electron, to become an cation. Because electronegativity decreases as you go down any Group, the elements can release their electron more easily as you go down a Group. Group VII elements react by stealing an electron from water, to become an anion. The more electronegative any element, the easier it is for it to steal an electron. Because electronegativity decreases as you go down any Group, the Group VII elements become less able to steal away an electron as you go down Group VII. |
No, there's no such thing as a semimetal bond.
Semimetals are elements that sometimes have some of the properties
of a metal (depending on other factors like molecular arrangement
and "contaminants").
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Some folks might call this a "trick question" but it's really meant to make you think about the shells.
The M-shell is capable of holding 18 electrons but it only does that when it is forced to do so by the larger outer shell. In the third Period the M-shell is the outer shell and it will only hold eight electrons (2 in an s orbital and 6 in the three p orbitals). That seems OK. But in the fourth Period the M-shell can be forced to store electrons in its d orbitals. That opens up a "basement" for storing an additional 10 electrons! So, in the fourth Period elements the outer shell (a N-shell) can force the M-shell to hold up to 18 electrons (2 in an s orbital, 6 in the three p orbitals and 10 in the five d orbitals). That's exactly what the first set of transition elements do. They (the transition elements in the fourth Period ) force the "basement" shell (the M-shell) to use its d orbitals.
The N-shell is larger than the M-shell and it can be forced to hold extra electrons in its "basement" by the shells above it. An element in the fifth Period will use its O-shell as an outer shell forcing the N-shell to use its d orbitals. That's exactly what the second set of transition elements do. They (the transition elements in the fifth Period) force the "basement" shell (the N-shell) to use its d orbitals. But that's not the end of the story.
An element in the sixth Period can use its P-shell to force the N-shell to use its f orbitals! That's exactly what the first set of inner transition elements do. They (the inner transition elements in the sixth Period ) force the "deep-basement" shell (the N-shell) to use its f orbitals. You will recall that there are 7 f orbitals so all together they are capable of holding 14 more electrons.
What of the O-shell?
You should be proud if you got that question right. Indeed, you should be proud if you understood that explanation the first time you read it. So read it again and focus on those "basement" problems. |
12C has 6 protons (that's why it's carbon, C), 6 neutrons (to
give it a total mass of 12) and has its electrons arranged as
2 (in the K-shell) and 4 (in the L-shell).
Carbon has 4 electrons in its outermost shell and it will try to get 4 more electrons to complete its outer shell, so it is a Group IV element. From other Tables we know that carbon strives to complete its outer shell by sharing electrons (thus forming covalent bonds) so carbon has a covalency of 4. |
40Ar has 18 protons, 22 neutrons and the electrons are arranged
as 2 (K), 8 (L), and 8 (a complete M-shell).
Because argon has a complete outer shell it is a noble element and belongs in Group VIII. Also, because its outer shell is complete, argon will be inert (not form bonds). It is unwilling to lose or gain electrons. |
Calcium's electron structure is 2 (K), 8(L), 8(M), 2(N).
If you guessed that it was squeezed into the M shell, you are
right!
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This isn't as easy as it might first appear, but I hope you got the outer electrons right. By looking at the table you should have guessed that the outer shell is the P-shell and it has two electrons (in an s orbital).
137Ba has 56 protons (you read that from the Table), 81 neutrons (you calculated by subtracting 56 from 137) and its 56 electrons are arranged as 2 (K), 8 (L), 18 (M), 18 (N), 8 (O) and 2 (in the P-shell).
The N-shell can also have d orbitals. In fact, it can have f orbitals too (if it must). The N-shell of barium has full s, p and d orbitals for a total of 18 electrons. That leaves only ten more electrons to assign somewhere. At this point it would have been an easy mistake to stuff the 10 electrons into the N-shell's f orbitals. But maybe you remembered that f orbitals aren't used until the NEXT shell (two shells up) is full. So move up to the next shell, the O-shell. Two electrons go into the O-shell's s orbital and six electrons go into the O-shell's three p orbitals. That leaves you with two electrons to assign. Where do they go? Perhaps you thought they would now go into the "basement" of the N-shell's f orbitals. After all, there's plenty of room there. Or maybe you wanted them to go into O-shell's d orbitals. Neither of these possibilities are right. Working with these "basement" orbitals is always difficult, because you never really known when to use them. That's where a look at the Table helps. If you found barium (Ba) on the Table, you would see that it is in the same Group as beryllium (Be), magnesium (Mg) and calcium (Ca). You know that they all have two electrons in their outer shell. And you can see from the Table that the P-shell is used. That's the clue you need to help you. The last two electrons are not in the "basement" orbitals of lower shells, they are in the s orbital of the P-shell! That's why barium is in Group II.
That was a very difficult puzzle to solve. You should be proud if you got it right. I didn't get that one right the first time I saw it. Even the explanation is hard to understand! So read through this answer again and be sure you understand it before going onto the next problem. |
187Au has 79 protons, 108 neutrons and you might expect its electrons to be distributed as 2 (K), 8 (L), 18 (M), 32 (N), 19 (O) but that would have made the O-shell the last shell and you know that's not right because the table shows gold is in the 6th Period so it must have the outer electrons in its P-shell. Looking at the other elements in the same column as gold is of little help.
The K-shell and L-shell are easy. Two electrons go into the K-shell and 8 into the L-shell. That's 10 down and 69 electrons to go (!).
The O-shell takes 2 electrons into its s orbital and 6 more electrons into its three p orbitals. Easy! That leaves us with only 11 electrons to go. But where do they go?
So the electronic configuration for gold is
Phew! That was hard. Please don't be disappointed if you could not figure out the outer shell of gold. Unlike the previous question about barium, the Table offers very few clues to guide you. Besides, all that hard work describing how the electrons fill each shell and fight over positions in the various orbitals is really something that very few Alchemists worry about. Any Alchemist working with gold would simply look up the electronic structure of a gold atom in a special reference book. The transition (and inner transition) elements fill their shells
in a more complex manner than the "non-transition"
elements. This gives the transition (and inner transition) elements
their unusual properties.
Oh, by the way, the transition elements don't go into real Groups. Actually, some 20th century Alchemists have labeled the transition "columns" with Group names (like "1B" for the column containing gold, silver and copper). Unfortunately there is no global standard for naming the transition elements' "Groups" (in the 20th century, but in a few more years..?). |
At first the number of protons equals the number of neutrons.
But as the elements get larger they take on more than an
equal number of neutrons.
Note: unstable isotopes (radioisotopes) will be "off" the green line, having too few or too many neutrons to make the nucleus stable. So they release or convert nucleons to reach a stable point somewhere on the green line. Radioactivity is caused by this reshuffling of nucleons as the atom tries to reach the stable position (on the green line). |
I hope you got these symbols for those elements and their isotopes
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Let's start thinking about the shells by their first quantum number.
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The sequence of orbitals produced in this way will look a wee bit out of order (some principal quantum numbers won't be in order) but that is good because this series represents the orbitals in the order in which they are filled.
This is a very useful technique and you should try it in the next few questions.
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Here is the pattern I got.
But that's not the end of our problem. How are the electrons arranged in that last shell. Look carefully at the list and note all those with principal number 7. Those include 7s2 and 7p6 and no others. That means this atom has eight electrons in its outer shell. Indeed, it has a complete set of eight in its outer shell. You would have been right to think that this imaginary atom is a noble gas! It would have all the properties you would expect of a Group VIII element. |
Here's a long string of the orbital series that you worked out earlier.
(I've only included the part of the series you need for these problems.)
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6, 6s2, 4f14, 5d10
Oxygen's 8 electrons go 1s2, 2s2, 2p4.
Carbon's 6 electrons go 1s2, 2s2, 2p2.
Argon's 18 electrons go 1s2, 2s2, 2p6, 3s2, 3p6.
Barium's 56 electrons took some work but hopefully you decided they are arranged as
And finally there was gold! You had to extend the series a bit further but I hope you found
Alchemists make up excuses for the behavior of the transition elements that misbehave. Actually, these aren't excuses, they are probable causes. In the case of gold the excuse (I mean cause) is like this. For a brief period of time the two outer shells are exactly as you have figured out above
You have to pay attention when adding up these orbitals. Notice that barium and gold both end up with a P-shell(6s), but gold has many more electrons stored in its "basement" shells. It is very important to follow the series when assigning the electrons to the orbitals and it is equally important to add the shells up correctly. |
Here's a string of the orbital energies we will need and I’ve listed them as all filled.
Nickel has 28 electrons so I can fill them as
Copper has 29 electrons so I can fill them as
Zinc has 30 electrons so I can fill them as
There's something special happening here. I checked the book and see that we have found the correct electron configuration for nickel and zinc but copper's M-shell and N-shell are wrong! The book says they should be
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Argon (Ar) is the last noble element before reaching the fourth Period where nickel, cooper and zinc reside.
Nickel is argon with the addition of a 3d8 and a 4s2 so it is shorter to write it as [Ar]3d8, 4s2. Notice that I am adding those extra electrons to Argon's M-shell (the 8 electrons in the 3d orbitals) and two electrons into the new shell that argon does not have, the N-shell (s orbital electrons).
Copper's correct configuration, using argon as a platform to build on, is [Ar]3d10, 4s1.
Zinc is [Ar]3d10, 4s2. Notice that you still have to start with the same old long series of electron orbitals but this method makes it easier to write the final answer and makes it more obvious how each atom differs from the others. Be sure to write this shorthand method into your notes. |
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There is only one type of s orbital and it is the orbital in the outer shell of elements in Groups I and II. It can hold only two electrons (because that is the maximum that any single orbital can hold). P orbitals help shape the atoms of the "post-transition" elements (Groups III to VIII) and there are three types (x, y and z) which all together can hold a total of 6 electrons. They are part of the outer shell of the post-transition elements (which they share with the two s orbital electrons). Indeed, s and p orbitals are the make up the outer shell of all atoms! the next two orbitals are “basement” orbitals. D orbitals are found in the "almost" outer shell of transition metals. That is, the transition elements place electrons into the d orbitals of their second outermost shell so d orbital electrons are never in the outer shell. (Not and warning: there is one element that breaks this "rule" so it isn't a rule it is a trend.)
The inner transition metals have f orbitals in their third outermost shell but never as their outer shell.
Note: All the transition elements (with one exception) have s orbitals in their outermost shell and d orbitals in the shell beneath. All the inner transition elements have s orbitals as their outer shell and (with one exception) f orbitals buried in their third outermost shell with some inner transition elements holding d orbital electrons in the second outermost shell. (Yes, they are very confusing atoms and they are a real headache to most Alchemists trying to make sense of these exceptions.) |