AQA – INORGANIC CHEMISTRY: Atomic structure – Flashcards

an atom fills up the sub shells in order of increasing energy :(NOTE: 4s is lower in energy than 3d and so will get filled in first) 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p - s orbitals hold up to 2 electrons - p ortbitals hold up to 6 electrons - d orbitals hold up to 10 electrons e.g. Sodium (Na) 1s2 2s2 2p6 3s1 BOX NOTATION: orbitals can also be represented by box notation. each box holds up to two electrons. each electron has an opposite spin represented by an up or down arrow remember: electrons prefer to occupy orbitals by themselves so only pair up when no empty orbitals of the same energy level are available - singly fill up all orbitals first and then add the remaining electron(s) with the opposite spin

the periodic table is split into blocks

When a positive ion is formed electrons are lost from the outermost shell Mg = 1s2 2s2 2p6 3s2 Mg2+ = 1s2 2s2 2p6 When a negative ion is formed electrons are gained O is 1s2 2s2 2p4 O2- is 1s2 2s2 2p6

4s is filled in before 3d and so when ions are formed 4s loses the electrons first Exceptions to the rule: Copper and Chromium are exceptions to this rule: Cr = [Ar] 4s2 3d4 Cu = [Ar] 4s2 3d9 The actual electron configurations are: Cr = [Ar] 4s1 3d5 Cu = [Ar] 4s1 3d10 This is to stabilise the structure - chromium would have had four boxes with paired electrons and one box with one unpaired electron, in the 3d orbital. To stabilise this, an electron from the 4s orbital is removed and added to the 3d orbital. that way all five boxes in the 3d orbital will have pairs of electrons spinning in opposite directions.

The enthalpy change when one mole of electrons is removed from one mole of atoms in the gaseous state to form one mole of singly charged positive ions in the gaseous state M (g) -> M+ (g) + e- M+ (g) -> M2+ (g) + e-

- the first ionisation energy down the group decreases - there is more shielding by increased number of energy levels between the nucleus and the outer electrons - increased shielding means that the atomic radius has increased, increasing the distance between the nucleus and outer most electrons - weaker attraction - less energy needed to remove the outer electron NOTE: as the number of protons in the nucleus increases going down Group 2 you might expect the first ionisation energy to increase because the nuclear charge increases. This doesn't happen because the factors described above have a greater influence on the value of the first ionisation energy.

- generally increases across a period - More protons in the nucleus so the nuclear charge increases - force of attraction between the nucleus and outer electrons increases (this pulls the electrons in closer to the nucleus reducing the atomic radius) - no addition of new levels to shield outer electrons because each successive electron enters the same energy level - more energy needed to remove the outer electron(s) as there is a stronger attraction between nucleus and electrons

Magnesium: 1s2 2s2 2p6 3s2 Aluminium: 1s2 2s2 2p6 3s2 3p1 The outer electron in aluminium is in a p sub-level. This is higher in energy than the outer electron in magnesium, which is in an s sub-level, so less energy is needed to remove it.

Phosphorus: 1s2 2s2 2p6 3s2 3p3 Sulphur: 1s2 2s2 2p6 3s2 3p4 It's not immediately obvious what's going on until we look at the arrangements of the electrons: The 3p electrons in phosphorus are all unpaired. In sulphur, two of the 3p electrons are paired. There is some repulsion between paired electrons in the same sub-level. This reduces the force of their attraction to the nucleus, so less energy is needed to remove one of these paired electrons than is needed to remove an unpaired electron from phosphorus.

1. The nuclear charge: the more protons in the nucleus the greater the charge so the higher the ionisation energy 2. The atomic radius: the bigger the atomic radius the further the outer electrons are from the nucleus, decreasing the attraction which decreases the ionisation energy 3. Shielding: if there is more shielding between the nucleus and the outer most electrons then the ionisation energy will drop because their is less attraction between the nucleus and outer most electrons

- used to determine all the isotopes present in a sample of an element and to therefore identify elements - needs to be under a vacuum otherwise air particles would ionise and register on the detector

Many ways to ionise a sample: 1. Electron impact vaporised sample is injected at low pressure electron gun fires high energy electrons at the sample knocks out an outer electron to form a positive ion with different charges e.g. Ti -> Ti+ + e- NOTE: electron impact is used for elements and substances with low formula mass because it can cause larger organic molecules to fragment 2. Electro Spray ionisation sample is dissolved in polar solvent injected through fine hypodermic needle to give fine mist or aerosol the top of the needle has high voltage at the top of the needle the sample molecule gains a proton from the solvent solvent formed evaporates away while the ions move towards a negative plate NOTE: used for lager organic molecules

- positive ions are accelerated by an electric field towards neg plate - given that all particles have the same kinetic energy, the velocity of each particle depends on the mass - lighter particles have a faster velocity and heavier particles have a slower velocity

- ions reach the detector and generate a small current - current fed into a computer for analysis - current is produced by electrons transferring from the detector to the positive ions - size of the current is proportional to the abundance of the species - for each isotope the mass spectrometer can measure a m/z and an abundance - sometimes two electrons may be removed from a particle forming a 2+ ion

the signal from the detector is passed to a computer where a mass spectrum is generated

- the relative atomic mass quoted on the periodic table is a weighted average of all the isotopes - R.AM = sum of (isotopic mass x %abundance) / 100