Basic ideas about atoms

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  • Basic Ideas about atoms    (1)
    • Atomic Structure
      • Nucleus of an atom contains protons and neutrons and is where the mass (most concentration of mass) sits. The overall charge is positive.
        • The Shell - or even orbit, or energy levels, contain electrons, a negative charge and a mass of 1/2000 which - obviously is negligible.
          • The overall charge of an atom is neutral as it has the same number of positive and negative atoms.
        • Energy levels of the nucleus have a negative charge.
      • Neutron; has a relative mass of 1 and a relative charge of 0.
        • Proton; has the relative mass of 1 and the relative charge of +1
          • Electron; has the relative mass of 1/2000 and the relative charge of -1.
            • The importance of the world 'relative' = the mass of an atom compared to a different atom.
    • Atomic Number and Mass Number
      • Mass Number - Number of particles present that have a significant mass - i.e. the neutrons & protons. It tells us the number of particles present in the nucleus (NUCLEONS)
        • Atomic Number - identifies an element. No two elements share the same atmoic number. It tells us the position in the periodic table, the number of protons in the nucleus (equal to the number of electrons in the energy levels - of course!)
          • Why shouldn't we define atomic number in terms of electrons?                               Well obviously because element can gain and loose electrons to form ions; why else?
    • Ions
      • An ion is an atom/group of atoms that has gained a charge.
        • A positive ion is called a CATION (like a cat) while a negative ion is called an anion.
          • Positive ions form through the loss of electrons while negative ions form through electron gain.
    • Isotopes
      • Atoms of the same element that has a different mass number. Due to this it has a different number of neutrons.
        • E.g. - there are two isotopes of chlorine, 35Cl and 37Cl. State the similarities and differences between atoms of these isotopes IN TERMS OF SUB-ATOMIC PARTICLES, ATOMIC NUMBERS AND MASS NUMBERS.
          • The atomic number remains the same, unlike the mass which changes.
            • 35Cl -   Protons  = 17   Electrons = 17  Neutrons = 18
            • 37Cl -        Protons = 17   Electrons = 17  Neutrons = 20
    • Nuclear Radiation
      • Radiation =          alpha       Nature = Highly ionising (helium nucleus)           Penetrating ability = weak (it doesn't go through paper)
        • Radiation=            Beta           Nature = Mildly ionising (high energy/speed electron)     Penetrating ability = mild (doesn't go through thin piece of aluminium)
          • Radiation =  Gamma         Nature  = weak ionising (electromagnetic wave) Penetrating ability = strong as it goes through a thick layer of lead.
            • As it is neutral it is not affected by the positive nor negative ends of the electric filed and follows a straight path.
          • As it is a high speed electron, beta has a negative charge so it is attracted to the positive end of the electric field.
        • As it is a positively charged nucleus, alpha is attracted to the negative end of the electric field.
    • Radioactive Decay
      • Alpha decay
        • 2 protons and 2 neutrons are emitted.    Atomic Number = two less than original  -2      Mass Number = Four less -4
          • Position of daughter isotope moves two to the left in the periodic table.
      • Beta Decay
        • A neutron in the nucleus becomes changed into a proton and an electron. The electron is then       emitted.     Mass number = remains the same              Atomic Number = + 1.
          • Position of the daughter isotope moves one to the right on the periodic table.
    • Half-Life
      • The time taken for the radioactivity of an isotope to fall to half its original value.
        • E.G. Potassium 40 is a radioactive isotope that decays by beta-emission and has a half-life of 1.25 x 10 (x9) years.
          • 1 ---->1/2 ----> 1/4 ---> 1/8      3 half lives      3 x 1.25x10(9)      3.75x10(9)
    • Electronic Structure
      • Instead of shell/orbits we now call them energy levels. Energy levels are numbered: 1 being closest to the nucleus, and containing electrons with the lowest energy. These numbers are called the Principle Quantum Numbers.
        • At each energy level, there are regions of space in which there is a high probability of finding an electron. These regions of space are known as orbitals.
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          • Any orbital can hold a maximum
    • Ionisation Energies
      • The molar FIRST ioniation energy of an element is the energy required to remove one mole of outermost electrons from one mole of its gaseous atoms.
        • If the ionisations are carried out under standard conditions (298K, 1atm) then these values are called the 'STANDARD MOLAR IONISATION ENERGIES'
          • The amount of energy needed to remove the outer electrons will vary from element to element
            • As you go across a period the ionisation energy increases.
            • As you go down a group, the ionisation energy decreases.
      • Variation in ionisation energies
        • The greater the attraction between nucleus and electron, the greater the ionisation energy.
          • The lower the attraction between the nucleus and the electron, the lower the ionisation energy.
            • 1. The attraction between nucleus and electron is influenced by two factors:
              • a. Nuclear charge - Nuclei have positive charges. A nucleus with a greater numbver of protons (a greater atomic number) will have a greater nuclear charge. e.g. 18Ar has a greater nuclear charge (18+) than 11Na (11+)
              • 2. Distance between the nucleus and electron is influenced by two factors
                • Electrons that are further from the nucleus are further from its 'pull'. E.g. A 4s electron is further from the nucleus than a 2s electron.
                  • However, an atom with a 4s outer electrons will also have a greater nuclear charge than an atom with 2s outer electron. The outermost electrons are shielded by inner electrons from the pull of the nucleus. So, even though the nucleus may have a greater charge,its effect is lessened through shielding.
      • The molar SECOND ionisation energy of an element is the energy required to remove one mole of outermost electrons from one mole of its gaseous uni-positive
        • Closer to the nucleus - stronger attraction - less shielding


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