Unit 1- Module 1: Atoms and Reactions

The first module in the AS chemistry course involves a lot of calculations and equations, and a few definitions to remember. It's all a bit daunting! But, if you break it all down, its all easy to remember.

The Atom

The Basic Structure of the Atom, with all the basic information you need to know about relative masses and relative charges of protons, neutrons, and electrons.

The nuclear symbol helps you to find out the number of protons, neutrons, and electrons in an element. 

Ions have different numbers of protons and electrons, with negative ions having more electrons than protons, and positive ions having more protons than electrons. 

Key Definition- Isotope- A form of an element which has the same number of protons and electrons, but different number of neutrons. 

Below are two examples of isotopes. 

Atomic Models

The 'atomic models' section in the OCR textbook and revision guide goes in depth into several models of the atom and the history of them. So, to make it more digestible, I have broken it down to the most important bits you need to know. 

Ancient Greeks first thought that all matter was made from indivisible particles. 

In the 19th century, John Dalton described atoms as spheres, saying that different types of sphere made up different elements. He still considered atoms to be indivisible. 

In 1897, J J Thomson found that atoms weren't solid and indivisible. He showed that an atom must contain even smaller negatively charged particles, which he called electrons. The model he drew up was called the 'Plum Pudding Model'. 

In 1909, Ernest Rutherford tested the 'Plum Pudding Model' by carrying out the 'Gold Foil Experiment'. 

• He fired alpha particles (which have a positive charge) at an extremely thin sheet of gold foil. 
• He expected most of the alpha particles to be deflected slightly, as the positive alpha particles and the largely positively charged sphere (in the atoms in the foil) would repel each other, so the particles would deflect off. 
• However, most alpha particles passed straight through the foil, with only a small number deflecting off at more than 90 degrees.

Rutherford called this new idea 'the Nuclear Model of the Atom'. This showed a tiny, positively charged nucleus at the center of the atom, surrounded by a 'cloud' of negative electrons. As most of the atom's mass is concentrated at the center, this model showed that most of the atom is empty space. 

Moseley later discovered that the charge of the nucleus increased from one element to another in units of one, so this must be due to positively charged particles, protons. 

They also noticed that as you increase the charge, the mass didn't increase proportionately. Later, James Chadwick discovered that this must be due to the neutron.

Finally, Bohr investigated and discovered that in Rutherford's model, the electrons in a 'cloud' would be attracted and pulled into the nucleus of the atom, which would then cause the atom to collapse. So he proposed a new model with four main principles:

• Electrons exist only in fixed orbitals (shells), nowhere in between
• Each shell has a fixed energy
• When an electron moves between shells, electromagnetic radiation is emitted or absorbed
• Because the energy of shells is fixed, the radiation will have a fixed frequency.

Bohr's model explained why some elements (noble gases) are inert. This is because the shells of an atom can only hold fixed numbers of electrons. When an atom has full shells it is stable. 

Key definition- Relative Atomic Mass- the average mean mass of an atom of an element  on a scale where an atom of carbon-12 is exactly 12.  

Some isotopic abundances may be given to you in the form of a graph, in which case they may not be given in percentages. Don't panic, you still work it out in same way, but instead of dividing by 100, divide by the sum of the isotopic abundances. 

Key definition- Relative Isotopic Mass- the mass of an atom of an isotope of an element on a scale where an atom of carbon-12 is exactly 12.

Key definition- Relative Molecular Mass- the average mass of a molecule on a scale where one alom of carbon-12 is exactly 12. 

Example:

Molecular mass of C₂H₄  = (2x12) + (4x1) = 28

Key definition- Relative Formula Mass- the average mass of a formula unit on a scale where an atom of carbon-12 is exactly 12. 

The Mole

The mole is a measurement of amount of substance. One mole is roughly 6.02x10²³ (Avogadro constant). 

There are three main calculations that you need to know involving moles.

Acids, Bases and Salts
Key definition- Acid- a substance that releases H⁺ ions in aqueous solution (a proton donor)

When mixed with water, all acids release H⁺ ions. 

eg HCl_(g) + H₂O ==> H⁺_(aq) + Cl^- _(aq)

H⁺ ions are just protons. They always combine with H₂O to form a hydroxonium ion (H₃0⁺).

Key definition- Base- a substance that removes H⁺ ions from an aqueous solution (a proton acceptor)

eg OH^-_(aq) + H⁺_{(aq) ==>} H₂O_(l)

Acid molecules release their hydrogen ions so that other ions can replace them. When the H⁺ ions have been replaced by metal ions or ammonium ions, you get a salt. 

There are 5 reactions of acids to produce salts that you need to know. These are:

      _{Acid + Base ==> Salt + Water}

      _{Metal oxide + Acid ==> Salt + Water}

      _{Metal + Acid ==> Salt + Water}

      _{Metal Hydroxide + Acid ==> Salt + Water}         

      _{Metal Carbonate + Acid ==> Metal Salt + carbon Dioxide + Water}

Metal hydroxide is an alkali, which is a base dissolved in water. When in water, it releases OH^- ions. 

Anhydrous and Hydrated Salts

All solid salts are arranged in a lattice, some incorporating water. This is called water of crystallisation. A solid salt that contains water of crystallisation is hydrated, one that doesn't is anhydrous.