Atomic Structure and Periodic Table

Chemistry A Level Notes: Atomic Structure and the Periodic Table

Atomic Theory:

Modern atomic theory developed by John Dalton (school teacher born 1776)
First ideas were developed by Democritus although he called them ‘atomos’ and his ideas came 2400 years before Dalton. The idea of atomos was deceased as Democritus realised if we keep cutting something at some stage we will be unable to further cut this thing. This thus creates a tiny particle which is unbreakable.
Dalton used measurements of the masses in which elements combine to propose that…

All matter made of atoms
All atoms of a given element are identical
A compound is a combination of two or more elements in a specific ratio
Atoms cant be created nor destroyed
In a chemical reaction atoms rearrange to give the reactions products

Doubts formed as Dalton viewed the atom as a spherical object which could not be cut any further and was somewhat like a marble.
J.J. Thompson discovered when all metals are heated they give off tiny negatively charged particles which were known as electrons. He was able to measure their mass to charge ratio
Milikan then measured the mass to charge ratio and discovered the weight of an electron was nearly 1/2000 of a hydrogen atom and so called them electrons. He proposed a ball of positive charge with electrons embedded in it, this was called the plum pudding model.
Becquerel theorised that if negatively charged particles are present so too must be positively charged particles for an atom to maintain structure. Becquerels radiation discovery led to him (along with marie curie) understanding that alpha particles are helium atoms which have lost two electrons.
Geiger and marsden bombarded gold foil with alpha particles after the suggestion of Rutherford that there were multiple types of radiation and that perhaps an atom consists of a tiny nucleus with electrons orbiting it. As the foil was very thin and only a few atoms thick, according to existing theories the particles should pass straight through.
Their actual results showed that…

An atom must contain a small, central, positively charged nucleus
The diameter of this nucleus must be 1/10000 of the atom
The majority of the atom's mass is centered in the nucleus
Electrons orbit the nucleus

They were able to conclude this because when they carried out the experiment, most particles went through the foil while some were deflected at an angle and did not progress straight while some others were deflected, hence hinting at the solid, positively charged, nucleus and the fact that the majority of others went through alludes to the fact that the majority of the atom must be empty space.
Thus Rutherford's concluding theory was that...

The nucleus of an atom of a particular element contains a fixed number of positively charged protons. This number is called the atomic number
The nucleus also contains a number of neutrons but this number varies from one isotope to the other, neutrons are not charged.
The mass number of an isotope is the total number of protons and neutrons in the nucleus of an atom of that isotope
The number of electrons in a neutral atom of an element is equal to the number of protons in the nucleus.
Electrons orbit the nucleus

Table of Knowledge for Atomic Summary

Key Definitions:

Atomic Number - the number of protons in the nucleus of an atom of that element
Relative Atomic Mass (ram) - the weighted average mass of an atom of an element compared to 1/12 the mass of a carbon-12 atom
Relative Molecular Mass - the average mass of a molecule of that element or compound compared to 1/12 the mass of a carbon-12 atom
Mass Number - the sum of the number of protons and neutrons in the nucleus of an atom of that isotope
Relative isotopic mass - the mass of an atom of that isotope divided by 1/12 the mass of a carbon-12 atom
Isotope - atoms of the same element that have the same number of protons but a different number of neutrons and so have the same atomic number but a different mass number. (same chemical properties, different physical properties)

Mass Spectrometry:

The masses of atoms, molecules and fragments of molecules can be measured using a mass spectrum. This consists of 4 key stages

Ionisation
Acceleration
Deflection
Detection

Initially all air is removed from the spectrometer creating a vacuum
A sample element or compound, which must be in a gaseous form is then injected into the mass spectrometer and this sample is bombarded with a beam of high energy electrons from an electron gun (even a solid element/substance must be heated adequately to create enough gaseous particles for the spectrometry to work). The energy from the bombardment of electrons leads to the atoms/molecules forming positive ions. (A → A+ + e-). These positive ions formed then fragment into multiple particles forming a positive particle and a neutral particle (minimum can be more) such as in chlorine whereby Cl2 → Cl2+ → Cl2+ + Cl.
The positive ions formed are then accelerated using the accelerator plate using high electrical potential. The ions then pass through slits, and emerge as a parallel beam of ions, as all the ions have the same energy if they have the same mass they travel at the same speed.
The ions are now deflected by a powerful magnetic field. Ions of greater mass but equal charge are deflected less and ions with lower mass are deflected more. The angle at which the ions are deflected is based on the strength of the magnet and the mass to charge ratio of the ions.
The ions of lowest mass will be detected first, the magnetic field will gradually be increased and the ions of higher mass are then detected. The detector is linked to a computer which calculates the mass to charge ratio of each positive ion and its relative abundance compared to the most abundant particle, this is then converted into a percentage for each positive ion. This is then used to calculate isotopic mass as an average, and is displayed as a mass spectrum graph.

In Brief...

Bombardment of high energy electrons to create positive ions
Acceleration of these positive ions by an electrical potential
Deflection of positive ions by a magnetic field
Detection of positive ions followed by the signals displayed as a graph

Often a mass spectra graph will have multiple lines, this is because many elements have isotopes and so there will be a line for every possible combination. Chloride for example should have three lines as the Cl-35 and Cl-35 ions with a mass to charge ratio of 70, equally the Cl-35 and the Cl-37 isotopes with a mass to charge ratio of 72 and the Cl-37 and Cl-37isotopes with a mass to charge ratio of 74 with additional peaks at 35 and 37 for the individual isotopes not in a molecular form
There are many types of mass spectrometers which vary based on use and how they are used, these include Time of Flight (TOF) spectrometers, where they are all loaded with equal kinetic energy and the heavier particles move slower. Another included Electrospray Ionisation (ESI) whereby samples undergo soft ionisation methods so there is no fragmentation, and so it is used to identify complex molecules.
Mass Spectrometers have many Uses:

Detection of Drugs and their metabolites in urine and blood samples

Most drugs such as certain steroids have a unique molar mass. Urine samples once taken from an athlete or testee, is placed in a low-energy mass spectrometer and ionised using the ESI method, it is then analysed and if the mass to charge ratio peak is equivalent to that of the molecular ion of steroid or drug then the testee must have the drug/steroid present

In the Pharmaceutical Industry:

The identity of a new compound can be determined using mass to charge ratio values. After a pure sample is obtained, a solution of substances can be sprayed through a needle of high electrical potential and is analysed to gain its mass to charge ratio.

Carbon-14 Dating:

The relative amounts of the isotopes of carbon-12 and carbon-13 in a sample can be found using mass spectrometry, as carbon-14 is radioactive its activity halves every 5730 years and changes to nitrogen-14, this method can then be used in conjunction with other methods to date rocks etc

Evidence of Electronic Structure & Ionisation Energies:

The energy needed to remove an electron from each atom in one molecule gaseous atoms is known as the first ionisation energy
The product is one mole of gaseous ions with one positive charge
First ionisation energy is written as (Na as example) Na(g) → Na+(g) + e-
Ionisation energies are always endothermic as energy is taken in the reaction and so energy sign will be +
Scientists can determine ionisation energies using a spectroscope which studies the light given out by atoms when heated in a flame as seen in flame tests.
Heating atoms gives them energy and this leads to the electrons becoming in an excited state. Once electrons are in an excited state they may jump to higher energy levels. When an electron does jump to a higher energy level it absorbs energy (can be radiation) the electron which has jumped must eventually drop back down to a lower energy level as it does not have the capability to maintain motion in the higher energy level. When the energy level drops back down it emits the same energy (can be radiation) that it absorbed upon i ncreasing in energy level. Each line in the spectrum which is seen arises from the energy given out when the electrons drop to the lower energy level (ground state)
It is possible to calculate the energy required to remove electrons from ions with increasing charges (known as successive ionisation energies)
There is an ionisation energy for each electron and an atom has, i.e for sodium there will be 11
The successive ionisation energies will increase, as once you remove one electron the second electron becomes more difficult to remove
Using Sodium as an example, you can expect a large jump between first and second ionisation energies and a large jump between the ninth and tenth ionisation energies. This is as a result of sodium atoms having a single electron in their outer shell (which is furthest from the nucleus). The outermost electron is very easy to remove as it is shielded from the charge of the nucleus and is the furthest away so feels the charge minimally. There is then a period of 8 electrons in the same basic energy shell which are closer than the outermost and not as shielded so are harder to remove than the outermost electron. The reason for a gradual increase in ionisation energies despite the electrons being in the same shell is that, as electrons are removed the overall charge of the ion becomes more and more positive, this means the force felt by each electron is greater and so the energy required to overcome this stronger attraction increases, hence the ionisation energy gradually increases for a period of 8 electrons. Finally there are two electrons in the innermost shell, which means there is another huge increase in the ionisation energy, these two ionisation energies are generally the highest as they are extremely close to the nucleus the nuclear radius is thus very small, there is also no shielding anymore and the nuclear charge is an extremely positive meaning the charge felt by the two electrons is now extremely high.
The electronic configuration of every electron can be represented in an energy level diagram or as an electronic configuration line 1s2 2s2 2p6 3s2
Electrons are found paired, for example in electron configuration diagrams the boxes are seen with two arrows in them when there are an electron pari present

Periodicity:

Repeating patterns in a period and are repeating through all periods
Melting Temperatures in a Period

The melting temperature of an element is based on its structure and its type of bonding between the atoms. In metals for example the bonding between atoms is very strong, this means their melting point is very high typically. The more electrons each atom contributes and so the bond is stronger and thus the more shared electrons the higher the melting temperature
This leads to a rise in melting temperatures from group 1 to group 4 and as carbon and silicon have giant covalent bonds and so the melting temperature is very very high. With very high graph peaks.
The non-metal groups 5,6,7,0 form simple molecular and the intermolecular forces are very wea (k so are easily overcome and so the melting boiling temperatures are very low

First Ionisation Energies of the Elements

Ionisation energies are determined by 3 key atomic properties.

The Nuclear Charge, the size of the positive nuclear charge. As the positive charge increases the attraction of the outermost electron increases and this increases ionisation energy
Nuclear Radius, which is the distance of the outermost electron from the nucleus, as this distance increases the ionisation energy decreases, and vice versa
Shielding Effect, Electrons in inner shells exert a repelling force on electron in outer shells and so the effect of attraction is reduced, the attraction of the nucleus is also shielde by these electrons reducing the attractive force

This leads to an increase in ionisation energy from left to right across a periodic table as the greater nuclear charges have a greater nuclear pull and thus more energy is required. Nuclear radius also decreases as the nuclear radius is decreasing due to the increased amount of protons present in the nucleus thus the shells are pulled closer. (NOTE: ionisation energy decreases down a group) shielding here stays the same so has no overall affect across a period although affects the decrease in ionisation energy down a group

Nuclear Charge Increases

Nuclear Radius Decreases

Shielding Level Stays the same