Unit 5 quiz

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Unit 5 Lecture

Lecture

Clovis Community college

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The Periodic Table is one of the most important tools in science. There is a ton of information jam-packed into this amazing library that can fit comfortably on one sheet of paper!

History

In 1863, John Newlands, an English chemist, classified the 56 known elements into 11 groups based on their physical and chemical properties. He noticed that when the known elements were placed in order of their increasing atomic masses, patterns involving physical and chemical properties sometimes repeated at intervals of eight, which reminded him of musical octaves. It turned out that he actually was right about some elements, but his colleagues thought that it was a pretty off-the-wall concept. There were problems with Newlands’ Law of Octaves, not the least of which was that it didn’t work after element 20, Calcium. Still it was a pretty neat observation – the patterns that he saw were very intriguing.

In 1869, the Russian chemistry professor Dimitri Mendeleev developed the first “real” Periodic Table. In one important way, Mendeleev followed Newlands lead – he also arranged the known elements according to their increasing atomic masses. What he found was that most of the time elements with similar physical and chemical properties fell into the same vertical columns (what he called families of elements). He also did something pretty darn clever – as he arranged the elements as to their increasing atomic masses there would be a big jump between one known element’s atomic mass and the next known element’s atomic mass. Rather than just smooch the two known elements together, Mendeleevleft a blank spot on his chart. He believed that there would be an element, to be discovered later, that would go there. He then made predictions about the properties of the unknown element based upon its location on his chart – and darned if he wasn’t just about on the money when the element would be discovered! That was a real selling point of Mendeleev’s chart – there had to be something to it for such accurate predictions to be made.

Something about Mendeleev’s chart bothered a young English scientist who was a student of Ernest Rutherford named Henry Moseley – it is the part that I have in italics above: most of the time elements with similar properties fell into the same vertical columns (or families) on Mendeleev’s chart, but not always. What was that about? In 1913, Moseley proposed that the chart should be based upon atomic numbers rather than atomic masses. At that time, there wasn’t accurate atomic number data for many of the elements so to back up his contention, Moseley did a series of experiments in which used some real cutting edge technology, x-rays, to determine atomic numbers. Turned out he was correct, when the Periodic Table was arranged as to increasing atomic numbers the elements that didn’t fit in the right family in Mendeleev’s chart moved over to the correct family. From everything I have read about Moseley, I firmly believe that he would have been a giant in chemistry and physics had he lived a full life, but when World War I broke out, he turned down a lucrative job offer and instead enlisted in the Royal Engineers. Tragically, he was killed in action in Gallipoli in 1915.

Take a look at the modern Periodic Table – It’s arranged by increasing atomic numbers, isn’t it?

Be sure to find a good Periodic Table on the World Wide Web and print it off to use. There are a zillion out there. The one that I have here is not the best copy in the world I just have it provided as an example.

Here’s some sites to try – but if they aren’t working, just type “Periodic Table” in your search engine and you will see a bunch!

http://www.mpcfaculty.net/ron_rinehart/periodic.htm (Links to an external site.)

http://cuip.uchicago.edu/www4teach/97/crothe/periodictable.html (Links to an external site.)

http://periodic.lanl.gov/downloads/PeriodicWebSite.pdf (Links to an external site.)

http://www.chemicool.com/ (Links to an external site.)

http://www.webelements.com/ (Links to an external site.)

http://www.chemicalelements.com/ (Links to an external site.)

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The Modern Periodic Law

(Puts what we’ve been talking about in a nutshell)

The physical and chemical properties of elements are periodic (repeating) functions of their atomic number.

Arrangement of the Periodic Table

As you can see, there are a couple of ways of looking at the Periodic Table. One way is to look at the horizontal rows on the chart. They are called periods or series (remember the series of elements when we were talking about electron arrangement? Same thing!) Another way of looking at the chart is to look at the vertical columns on the chart. Those are called groups or families (just like old Dimitri called them).

Let’s start with the 8 major groups or families.

Please note: There are a few different ways to identify the groups of elements. There are the old and new IUPAC (International Union of Pure and Applied Chemistry) systems and the CAS (Chemical Abstract Service of the American Chemical Society) and so on – and it can get pretty confusing. I am going to keep it simple and start with what I will refer to as the eight major groups – but please be aware that there are several other systems out there!

Look at the Periodic Table.

See Hydrogen? Look right above it, there is a 1 above the box that holds it – hydrogen begins major Group I. (Don’t let that little gap between hydrogen and lithium on the table above throw you – hydrogen, since it’s the simplest element in the universe, has some quirky properties – but it’s right where it should be, the top of Group I).

See that 2 above Beryllium? It’s the first element in major Group II
Now hop over that valley of the Transition Metals.
See that 3 above Boron? It’s the first element in major Group III
See that 4 above Carbon? It’s the first element in major Group IV
See that 5 above Nitrogen? It’s the first element in major Group V
See that 6 above Oxygen? It’s the first element in major Group VI
See that 7 above Fluorine? It’s the first element in major Group VII
See that 8 above Helium? It’s the first element in major Group VIII

Group I, a.k.a. the Sodium Family

With the exception of hydrogen, the members of this group are very active metals – these elements are so active that they are not found in a pure state in nature – they form compounds with other elements at the drop of a hat. Why are they metallic and so active? It all has to do with the number of electrons in their outermost shell – take a look: how many electrons are in the outer shells of hydrogen, lithium, sodium, and potassium? That’s right: only one electron in the outer shells of the members of Group I.

Group II, a.k.a. the Calcium Family

These elements are metals, but they are not quite as active as Group I metals; still, we can say they are active metals – they aren’t found in a pure state in nature either because they also form compounds with other elements very readily. Why are they like that? Those crazy outer shell electrons again! Look at Be, Mg, and Ca. Count up their outer shell electrons. How many did you come up with? That’s right: just two electrons in the outer shells of the members of Group II.

Group III, a.k.a. the Boron Family

Boron begins this group and it is considered a metalloid, where the other members of this group are given the name “poor metals” because they have lower melting points and are softer than their transition metal neighbors. But take a look at the members and count up their outer shell electrons: there are three electrons in the outer shells of the members of Group III.

Group IV, a.k.a. the Carbon Family

Carbon begins this group and is a non-metal, followed by silicon (a metalloid), and the rest of the members are poor metals. There are four electrons in the outer shells of the members of Group IV.

Group V, a.k.a. the Nitrogen Family

Nitrogen begins this group and is a non-metal, followed by another non-metal (phosphorus), then a couple of metalloids (arsenic and antimony), then a poor metal (bismuth). There are five electrons in the outer shells of the members of Group V.

Group VI, a.k.a. the Oxygen Family

Our old friend Oxygen begins this group and it’s a non-metal, as are sulfur and selenium, then tellurium and polonium (both metalloids) round out the group. There are six electrons in the outer shells of the members of Group VI.

Group VII, a.k.a. the Halogens

The name Halogen comes from a Greek word for “salt from the sea” and that’s a good name for this bunch of elements because several are found (not in a pure state, but in compounds) in ocean water: chlorides, bromides, iodides, fluorides. This is a very interesting group because it is the only one that has members in all three familiar states of matter – solid, liquid, and gas – at standard conditions. These elements are very active non-metals and the reason is found with the number of electrons in their outer shells. There are seven electrons in the outer shells of the members of Group VII.

Group VIII, a.k.a. the Noble Gas Family or the Inert Gases

Helium begins this group of very stable non-metals. They rarely react with other elements to form compounds – some never do. With the exception of helium (which has two), the members of this group have eight electrons in their outer shell. Another way of looking at it: helium has a full 1s orbital. It’s electrons are completely contained in the first energy level or K shell, and the rest of the members of this group have full s and p orbitals in their outer shells. This is why they are so stable and non-reactive – they are happy to be themselves!

Hey, that’s pretty cool! Whatever the number of the major group is, the members of that group will have that many electrons in their outer shells! (exceptfor He) That makes it easy to remember!

Other regions on the chart

Transition Metals

The transition metals are the 38 elements in the valley between major groups II and III. This region represents exactly what the name says: a transition from the active metals to the non-metals. The transition metals share a lot of the same characteristics as other metals but one big property that sets them apart from other metals is that their electrons that can combine with other elements to form compounds can be found in more than one shell that’s why some of them can have more than one charge as ions (we’ll get to that in the next unit – so file that away). Another way of saying that is to say that transition metals are able to put more than 8 electrons in the next-to-outermost shell – take a look at the Electron Configurations of the Elements chart from the last unit to check this out. Begin with scandium (atomic number 21). That’s the first element in the transition metals region. This is the point of the Electron configuration chart where the electrons begin backfilling into the 3d orbital even though there are electrons in the 4s orbital – see what I mean? This goes on through zinc (atomic number 30). After zinc, gallium (atomic number 31), everything is full and the 4p starts filling. The first transition metal in the 5th series is yttrium (atomic number 39), the last is cadmium (atomic number 48) – look what’s going on with the electron configurations: see how the electrons are backfilling into the 4d even though there are already electrons in the 5s? (OK, except for that weirdo element palladium!) So another way to think about the transition metals is to say they have electrons filling in the d orbital of the next to last shell.

Inner Transition Metals

Take a look at the two series of fourteen elements below the main body of the chart. These elements are called the inner transition metals.

These two rows, the Lanthanide series and Actinide series, have been separated from the main body of the Periodic Table to allow the elements that come after them (72 – 86 in the sixth series and 104 – 118 in the seventh series) to fall into the correct groups. Also, if they were included in the body of the table, it would be about twice as wide as it is now.

The Lanthanide series, elements 58 – 70 (the 14 elements beyond Lanthanum – that’s where the name comes from), are also known as the rare earth elements. Their 5d orbitals are filling even though they have electrons in the 6s.

The Actinide series, elements 90 – 103 (the 14 elements beyond Actinium – that’s where the name comes from), have their 6d orbitals filling even though they have electrons in the 7s.

No element beyond atomic number 92 is found in nature. They are all synthesized in labs.

Patterns on the Periodic Table

The Periodic Table has all kinds of repeating patterns running through it – that why It’s called the Periodic Table! Here are a couple:

  1. Atomic Radii (sizes of atoms)Let’s use the two ways of looking at the Periodic Table to see what happens with the sizes of elements/atoms.When you go down a group (take major group I, for example) each step down means you are adding a shell (hydrogen has just one shell, lithium has two, sodiumhas three, and so on). Because you are adding a shell with each step down, the atomic radii tends to increase (the sizes of atoms tend to get larger as you go down a group).When you go across a series (let’s use that second series that begins with lithium and ends with neon) something happens that you may not expect. As you move across a specific series, you are adding a proton (+ charge) to the nucleus with each step to the right – we know that because the atomic number is increasing by one with each step to the right – and we are adding an electron (- charge) to the same outer shell with each step to the right. In the case of the second series, Li has only one electron in the L shell, Be has two electrons in the L shell, B has three, C has four, etc. Since there is an extra positive charge being added to the nucleus and an extra negative charge being added to the same outer shell with each step to the right in the series, the natural attraction between positive and negative means the atomic radii tends to decrease (the sizes of atoms tend to get smaller) because the outer shell tends to be pulled closer and closer to the nucleus.
  2. Ionization Energy (the amount of energy needed to remove an electron from the outer shell of an atom)When you go down a group the sizes of atoms tend to get larger so the outer shells are further away from the nucleus. The further the outer shell electrons are from the nucleus, the less the attraction between negative and positive; the less attraction, the smaller the amount of energy you would have to have to remove an outer shell electron. I like to think of a satellite in Earth orbit as an analogy: the closer it is to the Earth, the more the Earth’s gravity pulls on it – the further away it is, the less it is pulled on by the Earth’s gravity. So, as you go down a group, ionization energy decreases.When you go across a series, the outer shells tend to be pulled closer to the nucleus, the attraction between negative and positive is stronger so ionization energy increases.

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