Topic 3: The Periodic Table and Atomic strucure

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The long form of the periodic table has its elements listed in order of atomic number(Z). This is the number of protons in an atom, and is unique to every atom. Listing atoms this way exposes elements with similar properties in a regular pattern. So when the list is arranged according the long-form, these elements are clumped together in a vertical column. This is called a group. The rows in the table are called periods.

 

A typical periodic table gives two pieces of information for each element. One is the atomic number, the other, the relative atomic mass (Aka molar mass) It has the symbol A. It is the mass of one mole of an element, taking into consideration the relative proportions of isotopes naturally occurring for the element. The masses of the isotopes are taken in relation to the Carbon-12 scale (Which is why it is called relative atomic mass). It is accurately determined using a mass spectrometer. It is calculated using this formula:

 

(Isotope RAM x percentage abundance) + (isotope x percentage abundance)...

 

Isotopes are elements with the same number of protons, but different number of neutrons.     

 

-Periodicity

 

When elements are placed in the long-form of the periodic table, we get the following patterns, or periodicity:

 

-          Molar volume (The volume of 1 mole of the substance) decreases towards group 4 then increases.

-          Density Increases to group 4 then decreases

-          Boiling points and melting points increase to group 4 then decrease.

 

Metals are found to the left of the table, metalloids near the centre (in groups 3 and 4) and non-metals to the right. The are classified according to their electrical conductivity, where metals are the best conductors, metalloids as semi-conductors and non-metals as the worst conductors.

 

-Ionisation energies and flame tests

 

The experiment:

 

1.      Dip a nichrome wire in concentrated HCl, and hold in a non-luminous Bunsen flame. If any flame colour appears, repeat procedure until no flame colour appears.

 

2.      When the nichrome wire is clean, dip in the concentrated HCl and then dip in the salt. Hold the nichrome wire with the salt on it in a non-luminous Bunsen flame, and observe and record the colour produced. View using a direct vision spectroscope to see different frequencies of light.

 

3.      To test for another salt, dip in HCl again, and hold in non-luminous Bunsen flame and repeat until no colour appears. Then repeat step 2.

 

We observe these colours for the following metal-salts:

 

Lithium: Scarlet

 

Calcium: Brick-Red

 

Strontium: Crimson

 

Barium: Apple Green

 

Sodium: Orange

 

Potassium: Lilac

 

Magnesium: No colour

 

Caesium: Blue

 

Rubidium: Red

 

Beryllium: No colour

 

-Why?

 

When an atom is energized (I.e. heat) electrons move from their lower shell to a higher shell, then collapse down to their lower shell again. This return to the lower shell releases energy in the form of radiation (i.e. a photon). A photon is a particle of light, and this is where the colour comes from. This is called excitation.

 

-Why the different colours then?

 

When viewed through a direct vision spectroscope, the flame tests produce definite lines, not a continuous spectrum. These lines form the line emission spectrum, which is unique for each element (can be used to identify elements). These lines appear only at certain frequencies. Because they appear at definite frequencies, then there must be a definite size for the jump between shells (distance between shells).  So an element with more shells will give more line emission spectra, because it will have more shells for the electrons to jump through, producing different frequencies.  Each element has a different number of electrons, or shells, therefore there can be large differences in the frequencies of photons produced.

 

When looking through a direct vision spectroscope at flame tests of Magnesium and Beryllium, we see line emission spectrum with frequencies higher than visible light. They produce Ultra-violet light.

 

When we look at the line emission spectra, let’s say for hydrogen, we notice that as we get closer to the ultra-violet side, the lines get closer (They converge, and approach their convergence limit). This tells us that the energy levels are getting closer the further we travel from the nucleus. The most energy unleashed in excitation can be achieved by the fall of an electron from a higher shell to the first shell (n=1).

 

-Ionisation

 

When an atom is supplied with enough energy, an electron can be removed. This leaves the particle as an ion. The energy needed to promote an electron outside the outermost shell is called the ionisation energy (Em) The first ionisation energy is the energy needed to remove one mole of electrons from one mole of atoms. These atoms are in a gaseous state.

M(g) --> M+ (g) + e- First Ionisation energy

 

M+ (g) --> M2+ (g) +e-  Second Ionisation energy

 

When plotted on a graph, the LOG for the ionisation energies can give us a startling pattern. From this graph we can deduce:

 

-          Ionization energies increase in general across a period (eg Li to Ne)

 

-          I.e decrease down a group (eg Noble gases)

 

I will explore reasons for each point, but first, I’ll need to explore atomic structure in more detail.

 

 

 

-Shells

 

Atoms are made of shells. We know that very well, but within every shell, are subshells. These subshells are made up of orbitals. Each orbital can hold a maximum of two electrons only. There are many types of subshell,  but we should only come across s, p, d and f. S has 1 orbital, P has 3 orbital, D has 5 orbitals, and F has 7 orbitals. Therefore, S can hold a maximum of 2 electrons, P a max of 6, D a max of 10 and F a max of 14. Electronic structure is written as follows:

 

Li- 1s2 2s1 (The number represents the shell, and the superscript how many electrons)

Be- 1s2 2s2

B- 1s2 2s2 2p1

C- 1s2 2s2  2p2

N- 1s2 2s2 2p3

O- 1s2 2s2 2p4

F- 1s2 2s2 2p5

Ne- 1s2 2s2 2p6

Electrons are more difficult to remove across a period for the following two reasons:

 

-          As we go across a period, nuclear attraction increases as there are more protons. This means it is more difficult to remove electrons. The atom also becomes smaller as electrons are attracted closer to the nucleus.

 

-          It takes more energy to remove an electron from a full or half full sub-shell because it is more stable

 

So, this can be cleared with the following example:

 

Li- 1s2 2s1

Be-1s2 2s2

 

Its takes more energy to remove an electron from beryllium because Beryllium has more protons, therefore more attraction, and beryullium also has 2 full subshells, compared to lithium with half a subshell.

 

But we notice that there are ‘blips’ at two stages in this period, between Be and B, and N and O.

 

Here is the electronic structure of Be and B:

Be- 1s2 2s2

B-1s2 2s2 2p1

 

It is easier to remove an electron from Boron because there is only one electron in the P subshell. This is less stable than a full subshell, which Beryllium has. Thus, less ionization energy is needed to remove the electron in the p- orbital of boron than of the s-orbital in beryllium.

 

The second blip features Nitrogen and Oxygen. This time, it is slightly more complicated:

 

N- 1s2 2s2 2p3

O- 1s2 2s2 2p4

 

Nitrogen has a half-full p-subshell. This is more stable than the p-subshell in oxygen, which is more than half full. The fourth electron joins another electron in a p-orbital, and because of repulsion forces between them, it is easier to remove than electrons in a half-full subshell.

 

-Down a group.

 

Looking at the chart, we notice a general decrease in ionization energy down a group. This can be explained as follows:

 

-          As you go down a group, more shells are added to the atom and therefore it is easier to remove outer electrons because they are further from the nucleus. This is because the attraction varies by the inverse square law.

 

-          Because of more shells, the inner electrons will shield the outer electrons from the nucleus’ attraction.

-Back to subshells.

 

Now, we have seen what subshells are, but we haven’t seen how they fill, and in what order.

 

When electrons are filled in the orbitals, they always occupy the empty orbital first. So when three electrons fill a P subshell, there is one electron in each orbital.  

 

From the above graph, we can see that s-orbitals are less energetic than p-orbitals etc. BUT there is an interesting exception. An s-orbital from a following energy level is less energetic than the d-orbitals in that subshell. That means, when we are filling the shells, the 4s subshell fills before the 3d subshell. This is our order of filling:

 

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10

 

Notice how 5s fills before 4d

 

E.g.

Calcium- 1s2 2s2 2p6 3s2 3p6 4s2

 

This order of filling has given the following groups these names:

 

1 &2 –S-block where 1 has an s-subshell with 1 electron in its outer shell, and 2 has an s-subshell with 2 electrons in its outer shell.

 

3-8- P-block etc

 

Transition elements- d-block

 

-Evidence for the presence of shells, and how to classify an element by use of ionization energies only

 

This is a graph of the ionization energies of sodium. We can tell the following from it:

-          The last two electrons are on a separate shell, because they are significantly more difficult to remove than the others. Therefore, they must be the electrons in the first quantum level.

-          The next 8 electrons are on another energy level, because they are significantly more difficult to remove than the first one, which suggests they are in their own subshell, in the second quantum level.

-          The last electron is in the third energy level.

 

From these findings we can state that there is one electron in the outer shell making sodium a group 1 metal. We can also see three shells , making it in period 3. this can make us identify it easily if we were ignorant of the number of electrons. This is also evidence for the presence of shells.

 

-Ions

 

I’m going to let the book explain the experiments related to this. As for ions and their sizes:

 

Negative ions in relation to corresponding atoms: They are bigger. These atoms have gained electrons. This imbalance off negative charge causes extra repulsion, which makes the ion bigger. The more electrons added, the bigger the repulsion, the bigger the ion. This means that ions of group 5 are bigger than ions of group 6 (as they are gaining 3 electrons as opposed to group 6 gaining only 2. Group 6 is bigger than group 7 for the same reasons.

 

Positive ions in relation to their corresponding atoms: They are smaller. There is a net positive charge as there is a loss of electrons. This net positive charge in the nucleus attracts the electrons more, bringing them closer to the nucleus. This shrinks the atom. The other reason is that when atom lose electrons to form ions, they also lose a shell, and this contributes to the shrinking in size.