Chemistry AS - Chapter 01 - Atomic Structure

Three fundamental particles:

• Protons - found in the nucleus
• Neutrons - found in the nucleus
• Electrons - surrounds the nucleus

As they are incredibly small, we use relative masses and relative charges:

• Proton:
  • Relative Mass: 1
  • Relative Charge: +1
• Neutron:
  • Relative Mass: 1
  • Relative Charge: 0
• Electron:
  • Relative Mass: 1/1840
  • Relative Charge: -1

In a neutral atom, the number of electrons must be the same as the number of protons as their charge is equal in size and opposite in sign.

Protons and Neutrons are held together in the nucleus by the strong nuclear force which is much stronger than electrostatic forces that hold electrons and protons together in the atom.

Strong Nuclear Force overcomes repulsion between protons in the nucleus and acts only over short distances.

Electrons are found in a series of levels (orbits/shells) which get further and further away from the nucleus.

Electrons are situated in shells surrounding the nucleus. First shell is closest to the nucleus. Electrons fill this shell first. Then second, then third etc.

General formula for the number of electrons in each shell =2n^2 where n is the shell number.

If we know the number of protons in an atom, we also know the number of electrons it has because the atom is neutral.

We can draw an electron diagram for any element.

We can also draw electron diagrams for ions as long as we know the number of electrons.

Sodium has 11 electrons. Its ion has 10. Therefore it has a positive charge. (as it has lost 1 electron.) (Na+)

An oxygen atom has 8 electrons but its ion has 10. So it has a negative charge of 2- (as it has gained two electrons.)

We can write electron diagrams shorthand:

• Write the number of electrons in each shell, starting with the inner-most shell, continuing outwards.
• Seperate each number with a comma.
• Examples:
  • Carbon - 2,4
  • Sulfur - 2,8,6
  • Na+ - 2,8

The number of protons in the nucleus is called the atomic number (or proton number). It's symbol is Z

The number of electrons is equal to the number of protons because atoms are electronically neutral.

The number of electrons in the outer shell of an atom determines the chemical properties of an element; how it reacts and what sort of element it is.

Atomic number Z = Number of Protons

Total number of protons and neutrons (total no. of nucleons) is called the mass number. It's symbol is A

Nucleons are responsible for all the mass the atoms have as electrons weigh virtually nothing.

Mass number, A = no. of protons + no. of neutrons

Every single atom of a particular element has the same no. of protons and therefore same no. of electrons. However, no. of neutrons may vary.

Atoms with the same number of protons but different number of neutrons are called isotopes.

Isotopes of the same element react in exactly the same way due to the same number of electrons being present in the outer shell.

Atoms of different isotopes of the same element vary in mass number because of the different number of neutrons in their nuclei.

Example is Carbon:

• Z number: 6
• Has three isotopes:
  • A numbers: 12,13,14
• All three will react in the same way due to same number of electrons in their outer shells.

The most useful instrument for accurate determination of relative atomic masses (R.A.Ms) is the Mass Spectrometer.

R.A.Ms are measured on a scale on which an atom of Carbon 12 is defined as exactly 12.

No other isotope has a R.A.M defined as a whole number because protons and neutrons do not have relative masses of exactly 1.

Mass Spectrometry determines the mass of separate atoms or molecules.

The instrument is kept in a high vacuum so the ions do not collide with air molecules which could stop them reaching the detector.

Sample is investigated in a gaseous state. If it is a solid, it should be vaporised first before being into the instrument.

Step 1: Ionisation.

• A beam of electrons from an "electron gun" knock out electrons from the atoms/molecules of the sample so they form positive ions. 
• Nearly all lose one electron to form ions with 1+ charge but 5% lose two electrons to form ions with 2+ charge.

Step 2: Acceleration.

• The positive ions are attracted towards negatively charged plates and are accelerated to a high speed. 
• The speed they reach depends on their mass.
• Lighter ions move faster than heavier ions. 
• Some of the ions pass through slits in the plates and they form into a beam.

Step 3: Deflection.

• The beam of ions move into a magnetic field which deflects the beam of ions into an arc of a circle. 
• Deflection depends on its mass to charge ratio (m/z). 
• 'z' means charge. 
• Heavier ions deflect less than lighter ions as they will resist the magnetic field strength more than lighter ions. 
• 2+ ions deflect twice as much as 1+ ions with the same mass. 
• Deflection also depends on magnetic field strength. Stronger the field, greater the deflection.

Step 4: Detection.

• The magnetic field is gradually increased so ions of increasing mass enter the detector one after another. Ions strike the detector, accept electrons, lost their charge and create a current which is proportional to its abundance.

• From the strength of the magnetic field at which an ion hits the detector, a computer works out the m/z value of the original ion.
• A read out called a mass spectrum is produced and is presented as a graph with relative abundance against m/z ratio.

Mass Spectrometer can detect different isotopes that make up an element. It detects individual ions so different isotopes can be detected due to the different masses of the ions.

Measuring R.A.Ms up to 5 decimal places is called high-resolution mass spectrometry. Work done to the nearest whole number is called low-resolution mass spectrometry.

To work out the average R.A.M from low resolution mass spectrometry:

• Take the relative abundance. Multiply it by its m/z ratio
• Do this for all readings on the mass spectrum.
• Add the values up, divide by the total abundance (usually 100)
• You get the average R.A.M

Electrons in different shells have differing amounts of energy. We can represent these on an energy level diagram.

We call shells "main energy levels". We label them as 1,2,3 etc. Each can hold a maximum number of electrons according to the formula 2n^2 where n is the shell number.

Apart from the first level, which has an s sub-shell, the main energy levels are divided into sub levels.

Sub-levels: s,p,d,f

Level 2 has an s and p sub level

Level 3 has an s, p and d sub level

We consider the electron to be a cloud of negative charge. Electron fills a volume in space called its atomic orbital.

Different atomic orbitals have different energies. Each orbital has a number that tells us the main energy level it corresponds to: 1,2,3 etc.

Orbitals have different shapes. These are sub-levels. Described as s, p, d and f.

Shapes represent a volume of space in which there is a 95% probability of finding an electron and also influences the shape of the molecule.

S orbitals hold 2 electrons.

P orbitals hold 2 each, but come in groups of 3. Therefore 6 electrons.

D orbitals hold 2 each, but come in groups of 5. Therefore 10 electrons.

Three rules for allocating electrons to atomic orbitals:

• Atomic orbitals of lower energy are filled first.
• Atomic orbitals of the same energy fill singly before pairing starts.
• No atomic orbital can hold more than 2 electrons.

Ionisation Energy is the energy required to remove one mole of electrons from a mole of atoms in the gaseous state. Measured in kJmol^-1

First electron needs the least energy to remove it because it is being removed from a neutral atom. Second electron requires more energy as it is being removed from a 1+ ion. Third electron requires even more energy as it is being removed from a 2+ion. We call these successive ionisation energies.

We can find the no. of electrons in each main level by looking at the jumps in successive ionisation energies. A large jump indicates the next electron is in a new shell.

Ionisation energies generally increase across a period due to increasing nuclear charge which means more energy is required to remove an electron.

However, the increase is not regular for Period 3.

From Mg to Al, the energy required decreases. This is because in Al, the 3p orbital has an electron in it and the 3p orbital is of a slightly higher energy than the 3s orbital. Therefore less energy is required to remove it.

There is a small drop from Phosphorus to Sulfur. This is because in each 3p orbital in Phosphorus, there is one electron. However in one of the 3p orbitals in Sulfur, there are two. The repulsion between these two make them easier to be removed, despite the increase in nuclear charge.

There is a general decrease in first ionisation energies as you move down a group.

This is because the outer electron is in a main level which gets further and further away from the nucleus as you progress down the group.

Although the nuclear charge increases, the actual positive charge "felt" by an electron in the outer shell is less than the full nuclear charge because of the inner electrons shielding the nuclear charge from the outer shell. The shielding increases as you go down a group.