Chemistry- Atomic structure and The Periodic Table
History of the atom
Timeline
Democritus 460-370 BC
- Tiny particles even in different shapes
- He called them atomos
- However Aristotle disagreed saying that all matter is made of water, fire, earth and wind
John Dalton 1766-1844
- In 1808 John Dalton saw that common things broke down into the same elements
- Different size and mass
J.J Thomson 1856-1940
- Atoms packed full of spheres of positive matter embedded with negatively charged electrons
Earnest Rutherford 1871-1957
- He did an experiment showing there is space for particles to pass through
- However some particles bounced back showing there was a nucleus
Niels Bohr 1885-1962
- This model shows the electrons orbiting the nucleus
Werner Heisenberg 1901-1976
- He challenged Bohr saying it was impossible to know the speed and the exact whereabouts of an electron
- Also known as the Plum Pudding model
- First speculation
- Found out that an atom can not be created or destroyed
- This was the discovery of the electron
- He won a nobel prize, however this model did not stay long
- He found this by using earlier works from Planck and Einstein
- Discovered electrons orbit nucleus at fixed energy levels
- Able to jump from one shell to another
Atoms- An atom is the smallest part of an element that can exist on its own
Elements- substance that contains only 1 type of atoms
Molecules- A molecule is where 2 or more atoms are chemically bonded together
Compound- Atoms of 2 or more different elements chemically bonded together
Particle
Mass
Charge
Location
Neutron
Proton
Electron
1
+1
Nucleus
1
0
Nucleus
very small
-1
orbiting around the Nucleus
The atom
The number of electrons = The number of protons
The charge of electrons = The charge of protons
Number of electrons in each shell
Shell 1 - 2
Shell 2 - 8
Shell 3 - 8
Shell 4 - 2
Number of Protons + Number of Neutrons = Mass number
Number of Protons = Atomic Number
Number of neutrons = mass number - atomic number
Isotopes- Atoms of the same element with the same number of electrons and protons but different number of neutrons
Ions
An ion = a charged particle
When atoms gain and lose electrons they develop a charge
Loose electrons = positively charged
Gain electrons = negatively charged
They are charged because the amount of electrons is not equal to the amount of protons
With positive ions, The charge = the group number
With negative ions, the charge = the group number - 8
Most types of metals form positive ions
Most types of non-metals form negative ions
Group 1 metals
The metals
Li - Lithium
Na - Sodium
K - Potassium
Rb - Rubidium
Cs - Caesium
Fr - Francium
Trends
- The further down the column the more reactive they get
- All are white crystalline solids
- The melting points decrease as you go down the column
- They are all colourless solutions
- They are all electrolyes
- They all have high melting points and high boiling points
All the metals in Group 1 have 1 electron on the outer shell meaning they all have 1+ charge, this means they are IONIC
- They all have 1 electron on the outer shell
Another name for them is ALKALI METALS
Water mixed with alkali metals causes an explosion
Li, Na, K
have to be stored in oil
they are soft and when exposed to air in a matter of seconds you can see it tarnishes showing a reaction
because of how reactive they are
Sodium + Water
Na + H2O = NaO + H2
It floats on the water and fizzes
Potassium + Water
K + H2O = KOH + H2
It fizzes and gives off a purple flame also the gas ignites on its own
Rubidium + Water
Rb + H2O = RbO + H2
It explodes giving off a cloud of gas this happens immediately
Caesium + Water
Cs + H2O = CsOH + H2
It immediately gives off a big explosion
Francium is even more reactive, it is dangerous
Group 7
The non-metals
F - Florine
Cl - Chlorine
Br - Bromine
I - Iodine
At - Astatine
Also known as the halogens
They all have 7 electrons in the outer shell forming a 1- ion when they react
Trends
- The melting and boiling points increase as you go down the column
- They all go around in diatomic pairs
They react with other non- metals and form covalent bonds
They react with metals to form ionic bonds
- They have all typical non-metal properies
Low melting points
Low boiling points
Brittle
Dull
They do not conduct electricity
Do not conduct heat
Fluorine - pale yellow - gas
Chlorine - yellow, green - gas
Bromine - reddish brown - liquid
Iodine - violet - solid
These are all types of structure covalent
Group 0
The non-metals
He - Helium
Ne - Neon
Ar - Argon
Kr - Krypton
Xe - Xenon
Rn - Radon
They are also known as Noble Gases
Because they have filled outer shells, there is no chemical reaction as there is no way to gain or lose electrons
Their electron structure is special as they all have filled outer shells
Helium - 2 - Gas - White, orange
Neon - 2,8 - Gas - Red, orange
Argon - 2,8,8 - Gas - Violet
Krypton - 2,8,8,8 - Gas - Grey
Xenon - 2,8,8,8,8 - Gas - Grey, blue
Radon - 2,8,8,8,8,8 - Gas - Colourless
Transition Metals
They are the large block of elements between group 2 and 3
Trends
- They have the highest melting points and boiling points
- They have higher densities
- They are much stronger and harder
- They are all metals
- They have metallic properties
They are much LESS REACTIVE
They have more uses than Group 1
Transition metals have special properties as well
- They form different charges e.g iron (II) ions, Fe 2+
- They form coloured compounds e.g copper (III) sulfate is blue
- They are useful catalysts e.g iron in the haber process and nickel in the hydrogenation of fats
Ionic Bonds - When group 1 metals react with non-metals
Covalent Bonds - A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms
Perfect paragraph to explain the reaction between Sodium and Chlorine
The outer electron from the sodium atom is transferred to the outer shell of the chlorine atom
To form a sodium ion (Na+) and a chlorine ion (Cl-)
Both ions have filled outer shells
As these ions have opposite charges so they attract each other
This is called an ionic bond