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🌡 GCSE Chemistry ⚗ - Coggle Diagram
🌡 GCSE Chemistry ⚗
C2
History Of The Atom
The original atomic theory was that atoms were solid spheres - each making up different chemical elements
In 1904, J Thomson concluded that there must, in fact, be something smaller within atoms. He theorised that these were negative electrons within the positive 'pudding' 
In 1911, the plum pudding model was disproved through the use of alpha particles by Ernest Rutherford. He noticed that, when he fired alpha radiation at gold, most of it was allowed through but some was reflected back. This was unexpected in the plum pudding model because it was thought that the positive charge was to spread out across the pudding to have an effect. The new model includes electrons orbiting around the nucleus - which was positive.
Finally, Bohr discovered that actually, if electrons were just floating around in clouds, they would collapse. Bohr proposed instead that electrons were actually locked within a fixed orbit around the nucleus
In the modern model, we tend to state that:
- Protons are heavy and positively charged
- Neutrons are heavy and have no charge
- Electrons have hardly any mass and are negatively charged
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Atoms, Ions & Isotopes
All atoms of a given element have the same amount of protons - they have to by definition as the atomic number is the number of protons.
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Ions are formed when atoms no longer have an equal charge - i.e. When they gain or lose electrons. This is written by a superscript notation of a + or -. This means that an iron iron with an extra electron would be Fe-
Isotopes on the other hand will still have the same neutral charge as they vary the amount of neutrons rather than electrons. This will change the mass number displayed as neutrons have a RAM of 1
The Periodic Table
History
The earliest known periodic table that we know of was Dmitri Mendeleev. He arranged the known elements as they knew them back then (around 50) but the key genius was that he left gaps. His table was arranged by relative atomic mass but with come elements switched over to ensure that elements in the same vertical groups had similar properties
The gaps that Mendeleev left allowed him to predict the properties of the elements he theorised were yet to be discovered. One notable example was that he predicted the properties of what he called ekasilicone or, as we know call it, germanium
Reactivity
Elements with the same number of electron in their outer most shell will often react in similar ways to one another - hence why elements in the same group behave like this
As a rule of thumb, metals will usually form positive ions when they react while non-metals will not usually
The ultimate 'goal' of atoms is to gain a full outer shell - at which point it will become stable. This can be achieved by losing, gaining or sharing electrons
Metals will form metallic or ionic bonds when they react while non-metals can form either covalent or ionic bonds
Metals
Transition metals are very typical metals - they share the properties you may expect of a metal and they can form more than one ion. They also have high melting points, are often colourful and can be used as an industrial cataylist in some applications
Alkali metals are the elements in group one and they share many properties. These include: Increasing reactivity, lower melting / boiling points, higher relative atomic mass and higher density along with a decrease in hardness-
These group one metals have some common reactions including the fact that they tarnish by reacting with water in the air
These elements are often stored in oil to prevent tarnishing however Rubidium and Caesium have to be stored in even more carefully controlled conditions to avoid a violent reaction
These metals, when reacted with water have the same equation. They react to form hydrogen and a metal hydroxide. They also react with chlorine to produce a salt
Group 7
The group 7 Halogens have the opposite trends to the group 1 elements. In the group 7, reactivity decreases going further down the group as these want to gain an extra electron and the further away from the nucleus this new electron will be - the less likely one is to attach.
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Group 7 elements engage in displacement reactions. This is where less reactive elements get 'pushed out' of a compound by a more reactive one. This can be used to find the reactivity of compounds
Noble gases are all of the group 0 elements. They are all inert as they already have a full outer shell and they are non flammable. They do, however, have increasing melting points
Ionic Bonding
When a metal and non-metal react together, they form an ionic bond. This is where one compound 'gives' some of its electrons away to the other. If possible, the elements will react to form a full outer shell
Ionic compounds have what is referred to as a regular lattice (like this)
As solid compounds, ionic materials cannot conduct electricity as they are arranged in a regular lattice like this however when they are melted or dissolved (which they do easily), there is a free electron available to carry the charges.
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Metallic Bonding
Metallic bonding creates a giant structure. In this model, electrons in the outer most shell of atoms are delocalised in what is often called a sea of delocalised electrons.
Substances bonded through metallic bonding have a number of properties are shared across them all. This includes: very high melting points, electrical and thermal conductivity. Metals are also very malleable and ductile. This is because, in pure metals, the atoms all line up with one another - thus leading to them being able to slide over one another. This is thus why harder metals can be created with alloys as these have some different sized atoms that do not line up exactly with one another.
Metals have their very own reactivity series. This goes (from most to least reactive):
- Potassium
- Sodium
- Calcium
- Magnesium
- Aluminium
- Zinc
- Iron
- Lead
- Hydrogen (comparison only)
- Copper
- Silver
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