Lesson 7 : Atomic Structures of Diamond and Graphite, with their properties.

Diamond

1. Diamond is a crystalline form of carbon.

2. It has a giant molecular structure

3. Diamond is made up of carbon atoms. Each carbon atom has a 2.4 electronic configuration and needs 4 more electrons to achieve a stable 2.8 electronic configuration of neon. A carbon forms four single covalent bonds with four other carbon atoms to form a giant molecular structure.

4. Each carbon atom is bonded to four other carbon atoms in a tetrahedral arrangement.

Properties of Diamond.

-Hard

Diamond is the hardest known substance.

Due to the strength of the bonding in all directions and its rigid lattice structure, the carbon atoms cannot slide over each other. This strength in bonding and structural rigidity results in its hardness.

-High melting and boiling points

The collective strength of all the covalent bonds within the diamond gives it a very high melting and boiling point as much force is required to overcome these strong electrostatic attractions between diamond molecules.

-Electrical conductivity

Diamond does not conduct electricity as there are no free valence electrons nor mobile ions since they are all used in bonding.

Uses of diamond

Diamonds are rare and precious so they are used as gemstones in jewellery.

Due to its hardness, diamonds are used as tips of drills and other cutting tools.

Graphite

1. Graphite is another crystalline form of carbon.

2. It has a giant molecular structure

3. Graphite has a layered structure, hence it is a good lubricant as there are weak van der waal's forces between the layers.

4. In a layer, each carbon atom is covalently bonded to three other carbon atoms in a hexagonal arrangement to form rings of regular hexagons. Each layer becomes a giant molecule.

5. As each carbon only uses 3 of its 4 valence electrons in bond formation, it still has 1 valence electron which becomes delocalised and mobile, hence it can conduct electricity.

6.The layers of hexagonal rings are held together by weak van der waal's forces of attraction and so the layers can slide over each other.

Properties of Graphite

Hardness.

Graphite is a soft substance as the different layers can slide over each other due to the weak van der waals forces of attraction between them.

High Melting and Boiling Points.

The collective strength of all the covalent bonds within the graph layers give it very high melting and boiling points.

Electrical conductivity.

Graphite conducts electricity parallel to its layered structure due to the free delocalised valence electrons.

Uses of Graphite.

Dry lubricant as it is soft and slippery

Inert electrodes for electrolisis

Lesson 6- The Periodic Table

PERIODIC TABLE - an organizational system for elements. Elements are arranged in rows going from right to left called Periods and columns going up and down called Families or Groups .
Elements in the same period have the same number of energy levels.

The period number is the same as the number of energy levels

Atoms get more massive as you move from left to right across the periodic table
The atomic radii decreases as you move left to right across the periodic table

Elements in the same family / group have similar properties because they have a similar electron arrangement.

Metals are on the left hand side of the table

Non-metals are on the right-hand side of the table.

Metalloids are between the metals and non-metals.

Group I or Alkali metals - Elements whose atoms have 1 outer-shell electron; they are very reactive
Group II or Alkaline Earth Metals - Elements whose atoms have 2 outer-shell electrons
Group III - Elements whose atoms have 3 outer-shell electrons
Group IV - Elements whose atoms have 4 outer-shell electrons
Group V - Elements whose atoms have 5 outer-shell electrons Group VI - Elements whose atoms have 6 outer-shell electrons
Group VII or Halogens - Elements whose atoms have 7 outer-shell electrons
Group 0, sometimes called group 8 or Noble Gases - Elements whose atoms have full outer shells so they are very unreactive.

Lesson 5: A brief highlight of Ionic Bonding, Covalent Bonding and Metallic Bonding (Triple Period)

What is an ionic bond?

An ionic bond is the force of attraction between the opposite charges of an ion. One element in an ionic bond loses electrons, and another element must gain the electrons. Some atoms lose electrons to make the outside energy levels become more stable. Atoms become more stable when their outer most energy level has 8 electrons. Pure ionic compounds usually are crystalline solids, liquids, or gases. Many ionic compounds are binary compounds.

This shows an atom losing an electron.

Ionic compounds usually have much higher melting and boiling points than covalent compounds.
Ionic Compounds:

Compound
Melting Point
NaCl
8010C
MgO
28520C
NaBr
7470C
LiF
840C

A covalent bond is formed when electrons are shared between atoms. Covalent bonds are between non-metals and non-metals or hydrogen and non-metals. They share electrons so that both of them can have a stable octet.

Metallic Bonding

Metals consist of a lattice of positive ions through which a cloud of electrons move. The electrons are the valency electrons of the metal, e.g. for sodium they are the outermost electron from each atom. The positive ions tend to repel one another, but are held together by the negatively charged electron cloud.

Because of this presence of mobile electrons, metals are able to conduct electricity

Lesson 4 : Ionisation Energy

Defining first ionisation energy

Definition
The first ionisation energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+.

This is more easily seen in symbol terms.
It is the energy needed to carry out this change per mole of X

Things to notice about the equation
The state symbols - (g) - are essential. When you are talking about ionisation energies, everything must be present in the gas state.

Ionisation energies are measured in kJ mol-1 (kilojoules per mole). They vary in size from 381 (which you would consider very low) up to 2370 (which is very high).

All elements have a first ionisation energy - even atoms which don't form positive ions in test tubes. The reason that helium (1st I.E. = 2370 kJ mol-1) doesn't normally form a positive ion is because of the huge amount of energy that would be needed to remove one of its electrons.
Patterns of first ionisation energies in the Periodic Table

The first 20 elements

First ionisation energy shows periodicity. That means that it varies in a repetitive way as you move through the Periodic Table. For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar.

These variations in first ionisation energy can all be explained in terms of the structures of the atoms involved.

Factors affecting the size of ionisation energy

Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionisation energy shows a high attraction between the electron and the nucleus.

The size of that attraction will be governed by:

The charge on the nucleus.
The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it.
The distance of the electron from the nucleus.
Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.

If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Between it and the nucleus there are the two layers of electrons in the first and second levels.

The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons.

The outer electron therefore only feels a net pull of approximately 1+ from the centre. This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.

Warning! Electrons don't, of course, "look in" towards the nucleus - and they don't "see" anything either! But there's no reason why you can't imagine it in these terms if it helps you to visualise what's happening. Just don't use these terms in an exam! You may get an examiner who is upset by this sort of loose language.

Whether the electron is on its own in an orbital or paired with another electron.
Two electrons in the same orbital experience a bit of repulsion from each other. This offsets the attraction of the nucleus, so that paired electrons are removed rather more easily than you might expect.

Explaining the pattern in the first few elements

Hydrogen has an electronic structure of 1s1. It is a very small atom, and the single electron is close to the nucleus and therefore strongly attracted. There are no electrons screening it from the nucleus and so the ionisation energy is high (1310 kJ mol-1).
Helium has a structure 1s2. The electron is being removed from the same orbital as in hydrogen's case. It is close to the nucleus and unscreened. The value of the ionisation energy (2370 kJ mol-1) is much higher than hydrogen, because the nucleus now has 2 protons attracting the electrons instead of 1.

Lithium is 1s22s1. Its outer electron is in the second energy level, much more distant from the nucleus. You might argue that that would be offset by the additional proton in the nucleus, but the electron doesn't feel the full pull of the nucleus - it is screened by the 1s2 electrons.
You can think of the electron as feeling a net 1+ pull from the centre (3 protons offset by the two 1s2 electrons).

If you compare lithium with hydrogen (instead of with helium), the hydrogen's electron also feels a 1+ pull from the nucleus, but the distance is much greater with lithium. Lithium's first ionisation energy drops to 519 kJ mol-1 whereas hydrogen's is 1310 kJ mol-1.
The patterns in periods 2 and 3

*The general trend is for ionisation energies to increase across a period

Lesson 3: History of models of atoms by deluded Chemists/Physists (Interrupted by fire drill, henceforth very short.)

John Dalton thought that an atom is a sphere of matter that is the same throughout.

J.J. Thomson discovered that all atoms contain electrons, which are tiny, negatively charged particles. Thomson proposed that an atom is a sphere of positive charge. The electrons are mixed uniformly in the sphere.

Rutherford updated the model of the atom. He hypothesized that almost all the mass and all the positive charge of an atom is concentrated in an extremely tiny nucleus at the center of the atom.

Bohr described the atom as a planetary arrangement: electrons orbiting the nucleus.

Lesson 2 : Atomic subshells (electronic configuration)

The distance from the nucleus that the electron spins is called its energy shell, energy level, or orbital.



Each energy level can only hold a certain amount of electrons


i. The first shell (K level) can hold 2 electrons


ii. The second shell (L level) can hold 8 electrons


iii. The third level (M level) can hold 18 electrons


iv. The fourth level (N level) can hold 32 electrons


v. The fifth level (O level) can hold 50 electrons


vi. The sixth level (P level) can hold 72 electrons


Each energy level is completely filled before electrons fill the next level.
The number of electrons in the outermost level are called valence electrons.








The energy sequence of the first 24 subshells is given in the following table. Each cell represents a subshell with n and given by its row and column indices, respectively. The number in the cell is the subshell's position in the sequence. Empty cells represent sublevels that do not exist.

The first period
Hydrogen has its only electron in the 1s orbital - 1s1, and at helium the first level is completely full - 1s2.


The second period
Now we need to start filling the second level, and hence start the second period. Lithium's electron goes into the 2s orbital because that has a lower energy than the 2p orbitals. Lithium has an electronic structure of 1s22s1. Beryllium adds a second electron to this same level - 1s22s2.
Now the 2p levels start to fill. These levels all have the same energy, and so the electrons go in singly at first.
B
1s22s22px1
C
1s22s22px12py1
N
1s22s22px12py12pz1
Note: The orbitals where something new is happening are shown in bold type. You wouldn't normally write them any differently from the other orbitals.
The next electrons to go in will have to pair up with those already there.
O
1s22s22px22py12pz1
F
1s22s22px22py22pz1

Writing the electronic structure of an element from hydrogen to krypton

Use the Periodic Table to find the atomic number, and hence number of electrons.
Fill up orbitals in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p - until you run out of electrons. The 3d is the awkward one - remember that specially. Fill p and d orbitals singly as far as possible before pairing electrons up.

Remember that chromium and copper have electronic structures which break the pattern in the first row of the d-block.
Writing the electronic structure of big s- or p-block elements

Note: We are deliberately excluding the d-block elements apart from the first row that we've already looked at in detail. The pattern of awkward structures isn't the same in the other rows. This is a problem for degree level.
First work out the number of outer electrons. This is quite likely all you will be asked to do anyway.
The number of outer electrons is the same as the group number. (The noble gases are a bit of a problem here, because they are normally called group 0 rather then group 8. Helium has 2 outer electrons; the rest have 8.) All elements in group 3, for example, have 3 electrons in their outer level. Fit these electrons into s and p orbitals as necessary. Which level orbitals? Count the periods in the Periodic Table (not forgetting the one with H and He in it).

Iodine is in group 7 and so has 7 outer electrons. It is in the fifth period and so its electrons will be in 5s and 5p orbitals. Iodine has the outer structure 5s25px25py25pz1.

What about the inner electrons if you need to work them out as well? The 1, 2 and 3 levels will all be full, and so will the 4s, 4p and 4d. The 4f levels don't fill until after anything you will be asked about at A'level. Just forget about them! That gives the full structure: 1s22s22p63s23p63d104s24p64d105s25px25py25pz1.

When you've finished, count all the electrons to make sure that they come to the same as the atomic number. Don't forget to make this check - it's easy to miss an orbital out when it gets this complicated.

Barium is in group 2 and so has 2 outer electrons. It is in the sixth period. Barium has the outer structure 6s2.

Including all the inner levels: 1s22s22p63s23p63d104s24p64d105s25p66s2.
It would be easy to include 5d10 as well by mistake, but the d level always fills after the next s level - so 5d fills after 6s just as 3d fills after 4s. As long as you counted the number of electrons you could easily spot this mistake because you would have 10 too many.

An atom's nucleus is held together by the strong nuclear force.
If the numbers of neutrons and protons are very different, the nucleus can become unstable and undergo radioactive decay.
Some nuclei decay by emitting an alpha particle.
Other nuclei decay by ejecting a beta particle.
Transmutation is a process in which one element changes into another through radioactive decay.
Half-life is a measure of the decay rate of a nucleus.
It is the time needed for one half of the mass of a sample of a radioactive isotope to decay.
Radioactive isotopes are used in medicine and for the study of the environment.
Tracer elements with short half-lives are followed through living systems to study certain things.

RECAP ON ATOMIC STRUCTURE.

Lesson 1 (Atomic Structure Basics & Isotopes)

First, let us enjoy this tutorial.






1. Matter is made up of atoms. Atoms themselves are made up of three sub-atomic particles : protons, neutrons and electrons

Particle/ Mass/ Charge

Proton 1 +1

Neutron 1 1 0

Electron 1/1840 -1

2. The protons and neutrons are located at the centre of the atom in the nucleus.

3. Electrons are arranged in shells around the nucleus. The 1st shell can hold a maximum of 2 electrons and the 2nd and 3rd shells can hold a maximum of 8 and 18 electrons respectively.

4.Every element has its own proton number, symbolised by Z, and the nucleon number, symbolised by A.

5. The proton number of an element is the number of protons in its atom. As an atom has to be electrically neutral, the number of electrons is the same as the number of protons.

6. The nucleon number of an element is the total number of protons and neutrons in its atoms.

Eg. The symbol for sodium (Na) is written as

nucleon number (A) --------23

------------------------------------Na

proton no. (Z) --------------11

Thus there are 11 protons, 11 electrons and 12 neutrons.

7. Isotopes are atoms of the same element with different numbers of neutrons.

For example, Cl-35 and Cl-37 are isotopes as they have the same number of protons. They are from the same element. However, they have a different number of neutrons. This difference in the number of neutrons can affect certain physical properties such as boiling point. It does not affect any chemical properties.

In the case of ions, the number of electrons will not be the same as the number of protons as they will lose or gain electrons during formation.

Eg.