1.

 Atomic Structure

The Atom

• Mass Number: tells us the number of protons and neutrons in the nucleus TOP

• Atomic (proton) number : tells us the number of protons in the nucleus BOTTOM

• Nucleus: most of the mass of an atom and very small, contains neutrons and protons

• Electrons: orbit the nucleus in shells and take most of the space of an atom 

Ions and Isotopes

• Ions have a different number of electrons and protons.

Negative ions 

E.g O²- have gained electrons to gain a full shell of electrons and enable to form a more stable ionic compound 

P= 8 charge +8

N = 8 charge 0

E = 10 charge -10

Total charge +2 

Positive ions 

Na+ have lost electrons to gain a full shell of electrons and enable to form a more stable ionic compound 

• Isotopes - elements with the same number of protons but different numbers of neutrons 

History of the atom 

 1803 - John Dalton - Atoms are spheres and each element is made form different spheres

1897 - JJ Thomson - Discovered the electron. The atom wasn’t solid and was made up of other particles. The plum pudding model was developed 

1909 - Ernest Rutherford - discovered the nucleus, also the nucleus was very small and +vely charged. Atom was maining empty space and made up a -ve cloud 

(gold leaf experiment - positive alpha particles fired at a thin gold leaf, Most went through the gold leaf (mainly empty space) Small number deflected back ( they hit a small positive nucleus)

1913 - Neils Bohr. discovered a problem in Rutherford's model. The cloud of electrons could collapse into the positive nucleus! He proposed electrons were fixed energy shells 

• The experimental proof - When EM radiation is absorbed, electrons move between shells. They EMit this radiation when electrons move down into lower energy shells

Atomic model today - Electrons don’t have the same energy in shells. We have subshells, this explains ionisation trends

Time of flight mass spectrometer 

1. Vapourisation - the sample is vapourised so it can travel through the TOFMS

2. Ionisation - the sample is pushed through a nozzle making a high pressure jet. A high voltage is passed through causing the loss of an electron. A gaseous positively charged sample is produced. Called electrospray ionisation

3. Acceleration - the positive ions are passed through an electric field. Particles with lower mass/charge m/z ratio will accelerate quicker  

4. Ion drift - particles travel through with a constant speed and kinetic energy. They drift through and particles with lower m/z ratios travel faster

5. detection - ions are detected as electrical current is made when particle hits the plate particles with lower m/z will reach the detector first as they travel faster

• Relative atomic mass Ar - the average mass of an atom of an element when measured on scale on which the mass of an atom of12C is exactly 12

• Relative Molecular mass (mr) - The average mass of a molecule when measured on a scale on which the mass of an atom of 12c is exactly 12 

• Relative isotopic mass - the mass of an atom of an isotope of an element measured on scale on which the mass of an atom is 12C is exactly 12

• Save my exam, chemrevise, allery chem

Mass Spectra - isotopes

 m/z is just the mass of an isotope divided by the charge. As most have just  +1 charge this is the same as the isotopic mass

The abundance is always shown on the y axis . it can be shown as a % or as a nominal value. If it is % all your isotopes must add to 100% 

From this we can work out the relative atomic mass

(Abundance A m/z A) +(Abundance B m/z B) / total abundance 

Mass spectra - molecules

(86) peak shows fragments of the original molecules. The last peak is the M + 1 peak or the molecular ion peak. This is same as the relative molecular mass of the molecule

Electron configuration 

S, 1 orbital can hold 2 electrons

P, has 3 orbitals can hole 2 * 3 = 6 electrons

D has 5 orbitals can hold 2*5 = 10 electrons
F has 7 orbitals can hold 2*7 =14 electrons

[1s, 2s,2p,3s,3p,4s,3d,4p,4d]

We fill orbitals singly first then pair up. This is due to electron repulsion

Electron configuration - ions with ions you just add or remove electrons from the highest energy level first first

Electron configuration - Transition metals

Chromium and copper behave differently 

An electron from the 4S orbital moves into the 3D orbital to create a more stable half full or full 3d sub-shell respectively 

Eg Cr is 1s2 2s2 2p6 3s2 3p6 3d5 4s1  

Transition metal ions behave differently too

When they become ions they lose their 4S electrons before their 3D electrons 

Ionisation 

• Ionisation energy is the minimum amount of energy required to removed one mole of electrons from one mole of atoms in the gaseous state

Na(g) -> Na+ (g) + e- 1st IE energy = 495.8kJmol-1

(ionisation requires energy so they are always an endothermic process and have a positive value)

Shielding 

The more electrons shell between the positive nucleus and negative electron that is being removed the less energy is required. There is a weaker attraction

Nuclear Charge 

The more protons in the nucleus, the bigger the attraction between nucleus and outer electrons. This means more energy required to remove this electron 

Atomic size 

The bigger the atom the further away the outer electrons are from the n nucleus. The attractive force between nucleus and outer electrons reduces - easier to remove electrons   

Successive Ionisation

The removal of more than 1 electron from the same atom is called successive ionization

Mg+(g)-> Mg2+ e- 2nd IE energy = +1450

Jump in energy as removing electrons from shell closer to nucleus

General increase in energy as removing an electron from an increasingly more positive ion

1st ionisation trends - groups 

Ionisation energy decreases as we go down a group 

• The atomic radius increases as we go down the group. Outer electrons further from the nucleus. Attractive forceis weaker. Energy required to remove an electron decreases

• Shielding increases as we go down. More shells between nucleus and outer shell. Attractove forcer is weaker. Energy required to remove an electron decreases

1st ionisation trends - periods

Ionisation energy increases as we go across a period 

• As we go across the period there is an increasing number of protons in the nucleus. This increases the nuclear attraction

• Shielding is similar and distance from nucleus marginally decreases 

• More energy required to remove an outer electron IE increases

• AL - decrease - the outer most electron in alimunium sits in a higher energy sub shell slightly further from the nucleus than the outer electron in Mg

•  S - decrease at sulfur is evidence for electron repulsion in an orbital, electrons repel each other so less energy is needed to remove an electron from an orbital with 2 in than a one