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Topic 1: Atomic Structure and the Periodic Table

Atomic Structure

  • Atoms consist of a nucleus surrounded by shells of electrons
  • The nucleus consists of protons and neutrons
  • The protons in the nucleus causes a positive charge
  • The electrons orbit the nucleus and have a negative charge

Atomic and Mass Numbers

  • The Atomic Number is the number of protons in an atom
  • The Mass Number is the number of protons and neutrons in an atom
  • The Mass Number displayed on the Periodic Table is the average mass number based on the relative abundancy of the different isotopes of the element

Isotopes and Ions

  • An Isotope is an element whose atomic number is the same but the mass number(number of neutrons) is different.
  • Isotopes only have different numbers of neutrons. The number of protons and electrons is the same
  • Ions are charged atoms which have a different number of electrons
  • Positive Ions have lost electrons
  • Negative Ions have gained electrons
  • Protons determine the element. Changing the number of protons means a new element is formed

Relative Masses

  • The Relative Atomic Mass of an element is the average mass of all the isotopes of an atom relative to the mass of 1/12 of the mass of Carbon-12
  • The Relative Isotopic Mass is the mass of an isotope of an element relative to the mass of 1/12 of the mass of Carbon-12
  • The Relative Molecular Mass is the average mass of all the atoms in a moleculerelative to the mass of 1/12 of the mass of Carbon-12
  • Carbon-12 is used because it is a stable and abundant isotope

Mass Spectrometry

  • Mass spectrometers can be used to determine all the isotopes present in a sample of an element and their percentage or relative abundancy.
  • The relative atomic mass quoted on the periodic table is a weighted average of all the isotopes of an element
  • The data collected from a mass spectrum can help to calculate a relative atomic mass:
    RAM = Σ (isotopic mass x % abundancy)
    100

    RAM =Σ (isotopic mass x relative abundancy)
    total relative abundancy
  • Chlorine has 2 isotopes: Cl-35 (75%) and Cl-37 (25%) so the RAM is calculated to be 35.5
  • On a mass spectrum however, there will be three peaks at 70, 72 and 74 due the presence of the different isotopes in each molecule (35-35; 35-37; 37-37)
    Mass spectrometers are used for:
  • Drug testing in sports to identify chemicals in the blood
  • Quality control in the pharmaceutical industry
  • Radioactive Carbon-13 Dating to determine ages of fossils or human remains
  • Testing rocks on different planets

Ionisation Energies

  • The first ionisation energy is the energy required when one mole of gaseous atoms forms one mole of gaseous ions with a 1+ charge
    X(g) –> X+(g) + e-
  • The second ionisation energy is the energy required when one mole of gaseous ions with a 1+ charge forms one mole of gaseous ions with a 2+ charge
    X+(g) –> X2+(g) + e-
    Ionisation energies are affected by:
  • The attraction of the nucleus (the more protons, the greater the attraction)
  • The distance of the valence electrons from the nucleus (the further from the nucleus, the weaker the attraction)
  • The shielding provided by the electrons (the more electrons infront of the valence electrons, the weaker the attraction)
  • Helium has the highest first ionisation energy because its nuclear attraction is strongest as the atomic radius is shorter than other atoms and there is no shielding
  • Successive ionisation energies are always larger because there is a greater positive:negative charge ratio as there are fewer electrons and the ionic radius is smaller

Ionisation Energies and Electronic Structures

  • The more electrons removed from an atom, the higher the ionisation energy as there are more protons to electrons so the nuclear attraction is greater for each subsequent electron removed
  • A larger jump in Succesive Ionisation energy can be seen when an electron is removed from a shell closer to the nucleus as there is less shielding and the distance is smaller
  • Nobel Gases have higher first ionisation energies as they have the most protons with the least amounts of shielding
  • A large drop in first ionisation energies occurs between the Nobel Gases and Group 1 metals as shielding increases weakening the nuclear attraction and causing electrons to be more easily lost
  • Small drops in first ionisation energies can be seen between Group 2 and 3elements (e.g. Be and B) as the electrons are removed from a different subshell
  • Small drops in first ionisation energies can be seen between Group 5 and 6elements (e.g. N and O) as the electrons start to pair up in a subshell

Graph of First Ionisation Energies

Electronic Structure of the Atom

  • Electrons are arranged into principle energy levels (shells) which are split into sub energy levels (subshells) and can contain a certain number of electrons.
  • Subshells are split into
    s (holds up to 2 electrons)
    p (holds up to 6 electrons)
    d (holds up to 10 electrons)
    f (holds up to 14 electrons)
  • Subshells are made up of orbitals which hold up to two electrons of opposite spin
  • Orbitals represent the mathematical propabilities of finding an electron at any point within certain spatial distributions around the nucleus
  • s subshells are spherical
  • p subshells have a dumbbell shape
  • The periodic table is split into block based on the subshell that the valence electron is found

Electronic Configurations

  • An atom fills up the subshells in order of increasing energy
  • The 3d subshell is higher in energy than the 3d subshell so gets filled after the 4s
  • 1s > 2s > 2p > 3s > 3p > 4s > 3d > 4p > 5s > 4d > 5p
  • An electronic configuration includes the main energy level, the name of type of sub-level and the number of electrons in the sub-level e.g. Oxygen = 1s2 2s2 2p4
  • Spin diagrams can be drawn to show the number of electrons and the spin of the electrons in the orbitals of an element.
    When filling up sublevels with several orbitals, the orbital are filled individuallybefore electrons start to pair up

 

Periodicity

  • Periodicity is the trend of chemical and physical properties across a period or down a group of elements
  • Atomic Radius decreases across a period as there are more protons so the nuclear charge is greater and pulls the electrons in more
    Atomic Radius increases down a group as more shells and increased shielding weaken the nuclear attraction on the valence electrons
  • First Ionisation Energy tends to increase across a period due to an increase in protons providing a stronger nuclear attractive force on the electrons
    First Ionisation Energy tends to decrease down a group because the increased number of shells results in shielding weakening the nuclear attractive force
  • Melting Points are based on bonding and intermolecular forces within the elements.
    For Na, Mg and Al – These elements have metallic bonding so the increase in outer electrons and a smaller ion results in a larger melting point
    For Si – Silicon is a macromolecular substance so requires a lot of energy to break the multiple covalent bonds in the structure
    For Cl2, S8, P4 and Ar – These elements exist as simple molecular substances that have weak Van der Waals Forces between them so less energy is required to break these Intermolecular Forces. The more electrons they have in a molecule, the higher the m.p.

 

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