Atomic Structure
- Atoms are the components that make up all elements.
- Atoms are made up of three types of sub-atomic particles – protons, neutrons, and electrons
- Protons and neutrons make up the nucleus, where most of the mass is concentrated. Electrons orbit the nucleus in shells
Particle | Relative Mass | Relative Charge |
Proton | 1 | +1 |
Neutron | 1 | 0 |
Electron | 1/1840 | -1 |
The Evolution of Atomic Structure Over Time
- The model of the atom has changed over time as new evidence has become available.
Dalton proposed that all atoms of one element are the same and are different from the atoms of another element. Atoms in his model were tiny and indivisible.
Thomson discovered the electron. He proposed the plum pudding model where negatively charged electrons move in a ‘sea’ of charge in a positively charged atom.
Rutherford found that most of the mass is concentrated in the positive nucleus, with negative electrons orbiting it. The positive and negative charges balance to make the atom neutral.
Bohr suggested that electrons orbit the nucleus on paths. Bohr’s planetary model provided an explanation for the difference in energy of electrons at different distances from the nucleus.
The current model is composed of protons, neutrons, and electrons. Protons and neutrons are found in the nucleus and are made up of smaller quarks, whereas electrons surround the central nucleus.
Mass Number & Isotopes
- Element, X
- Mass number, A, is the total number of protons and neutrons in the nucleus
- Atomic number, Z, is the number of protons. The number of positively charged protons is equal to the number of negatively charged electrons in an atom, making the atom neutrally charged
- Mass number = number of protons + number of neutrons
- Atomic number = number of protons = number of electrons
- Ions are formed by atoms losing or gaining electrons.
- A charge of x- means that the number of electrons in the ion is the atomic number + x
- A charge of x+ means that the number of electrons in the ion is the atomic number - x
- Isotopes are atoms with the same number of protons and different numbers of neutrons. Therefore, they have different mass numbers but the same atomic number.
- Isotopes of the same element have the same electronic configuration so react in the same way in chemical reactions but have slightly different physical properties.
Mass Spectrometry
- Mass spectrometry is a form of molecular chemical analysis that allows the masses of individual molecules or isotopes to be determined.
- Mass spectrometry can be used to provide structural information, identify an unknown compound, or determine the relative abundance of each isotope of an element.
- In time of flight mass spectrometry, the steps include:
- Ionisation- the sample is dissolved in a volatile solvent and ejected through a hollow needle. The needle is connected to a positive terminal of a high voltage supply. This produces tiny positively charged droplets.
- Acceleration- ions are accelerated towards a negatively charged plate to give all ions constant kinetic energy. So, the velocity of each ion will depend on its mass
- Ion drift- ions pass through a hole in the negative plate, forming a beam
- Detection- the positive ion picks up an electron which causes a current to flow. Flight times are recorded
- Data analysis - the signal from the detector passes to a computer which generates a mass spectrum
- The mass spectrum gives information about the relative abundance of isotopes on the y axis and about the relative isotopic mass on the x axis.
- The mass spectrum can be used to determine the relative atomic mass (Ar)
- A mass spectrum for a molecular sample shows the relative molecular mass on the x axis.
Electronic Configuration
- Electrons orbit the central nucleus in shells. Each shell can hold 2n^2 electrons, where n is the principal quantum number.
- Electron shells are made up of atomic orbitals, which are regions in space where electrons may be found.
- Each shell is composed of one or more orbitals and each orbital can hold one pair of electrons.
- There are four main types of orbitals: s-, p-, d-, and f-
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- Electrons have an intrinsic property (spin). For two electrons in the same orbital, the spin must be opposite to minimise the repulsion.
- Within each shell, orbitals that are of the same energy level are grouped together in sub-shells.
- There are 1 s-orbital, 3 p-orbitals, 5 d-orbitals and 7-p orbitals possible in each subshell.
- Sub-shells have different energy levels. Note that 4s is lower in energy than 3d, so 4s will fill first.
- Shells and sub-shells are filled with electrons according to a set of rules:
- Atomic orbitals with the same energy fill individually first before pairing
- Aufbau principle – the lowest available energy level is filled first
- No more than two electrons can fill an atomic orbital
- Electron configuration is written with n representing principal quantum number. X is the type of orbital and y is the number of electrons in the orbitals of the subshell e.g. potassium has 19 electrons and its electron configuration is written as 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1
Ionisation Energy
- Ionisation energy is a measure of the energy required to completely remove an electron from an atom of an element to form an ion.
- First ionisation energy is the energy required to remove one electron from each atom in one mole of the gaseous element to form one mole of gaseous 1+ ions. X (g)→ X+(g) + e-
- Successive ionisation energies apply to the removal of electrons after the first ionisation energy. The nth ionisation energy is:
- Successive ionisation energies provide evidence for the shell structure of atoms.
- Within each shell successive ionisation energies increase, as there is less electron repulsion
- Between shells, there are big jumps in ionisation energies, as the electric is removed from a shell closer to the nucleus
- Factors affecting ionisation energies:
- Atomic radii- The larger the atomic radius, the further away the outer electrons are held from the nucleus, and the smaller the nuclear attraction.
- Nuclear charge- The greater the nuclear charge, the greater the attractive force on the outer electrons.
- Shielding- electrons repel each other due to their negative charge. The greater the number of inner shells of electrons, the greater the repulsion of the outer shell of electrons.
- The greater the attraction, the harder it is to remove an electron. Therefore, the ionisation energy will be larger.
- Atomic radii show periodicity. Across a period, the radius decreases while down a group, the radius increases.
- Ionisation energy increases across a period, electrons are all added to the same shell resulting in greater attraction.
- Ionisation energy decreases across a group, as the number of shells increases, so does the atomic radius and shielding, reducing attraction.
