Atomic Structure_ Bonding and Periodicity

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					Atomic Structure, Bonding and
•   Atomic Structure
•   Amount of Substance
•   Bonding
•   Periodicity
            Atomic Structure
•   Fundamental Particles
•   Mass Number and Isotopes
•   Mass Spectrometer
•   Electron Arrangement
•   Trends Ionisation Energies
           Fundamental Particles
Atoms have a small central nucleus made up of protons and neutrons
around which there are electrons.

Electron                              Atomic     Relative     Relative
                                      Particle   Mass         Charge
               +                      neutron        1            0
                                      proton         1           +1
                                      electron   negligible      -1

• Atoms of the same            Mass number: this is the number of
  element always have the      protons and neutrons (A)
  same number of protons.
• Isotopes are atoms of the
  same element but with              4
  different number of
  neutrons. This gives rise
  to different mass
                                     2   He
• Relative abundance is the
  amount of each isotope as     Atomic number: this is the
  the percentage for that       number of protons (Z)
  element occurring on the
                  Mass Spectrometer
                                           A curved magnet
               An           Charged
 The                                       deflects the
               electron     plates                              The ions are detected
 sample is                                 positive ions. The
               gun          accelerate                          to produce a mass
 put in via                                lighter ions are
               ionises      positive                            spectrum
 an injector                               deflected more
               the atoms    ions

               Ionisation   acceleration    deflection          detection

•The mass spectrum gives the relative           Example: Copper
masses of each isotope and the                  69% of copper atoms have a
abundance of each isotope.                      relative mass of 63.
•The relative atomic mass is the                31% of copper atoms have a
weighted average mass of an atom of an          relative mass of 65.
element compared with 1/12 of the mass          The weighted average is calculated
of 12C.                                         as follows:
•We can calculate this from mass                (0.69 x 63) + (0.31 x 65) =63.62
      Arrangement of Electrons
• Energy Levels or Shells
   – The simplest model of electrons has them orbiting in shells
     around the nucleus. Each successive shell is further from the
     nucleus and has a greater energy.
• Sub Shells and Orbitals
   – This model can be further refined by the concept of sub shells
     and orbitals.
   – Sub shells are known by letters s, p, d, and f. The s sub shell can
     contain 2 electrons, p 6, d 10 and f 14.
   – Electrons occupy negative charge clouds called orbitals, each
     orbital can hold only 2 electrons. Each type of shell has a
     different type of orbital.
          Arrangement of Electrons
•   How we write electron configurations
     – Electrons fill the lowest energy level first this means it is generally easy to predict how the
        electrons will fill the orbitals (it gets more complicated with the transition metals).

     Element       Electron                           Element        Electron configuration
     H             1s
                                                      Ne             1s22s22p6
     He            1s2
                                                      Na             1s22s22p63s1
     Li            1s22s
                                                      Mg             1s22s22p63s2
     Be            1s22s2
                                                      Al             1s22s22p63s23p3
     B             1s22s22p1
                                                      Si             1s22s22p63s23p3
     C             1s22s22p2
                                                      P              1s22s22p63s23p3
     N             1s22s22p3
                                                      S              1s22s22p63s23p4
     O             1s22s22p4
                                                      Cl             1s22s22p63s23p5
     F             1s22s22p5
                                                      Ar             1s22s22p63s23p6
                     Trends in Ionisation Energies
                   The first ionisation energies of group 2                       First ionisation energies of period 3
                                   electrons                                                     elements

                   1000                                                         2000
 Energy (Kj/mol)

                                                              Energy (Kj/mol)
                   800                                                          1500
                   200                                                           500

                     0                                                             0
                          Be      Mg      Ca      Sr     Ba                            Na   Mg   Al   Si   P     S   Cl   Ar
                                   Group 2 elements                                              Period 3 elements

Going down a group in the periodic table
there are more filled energy levels between the
                                                              Going across a period of elements:
nucleus and the outer most electrons
                                                              there are more protons in each nucleus so the
these shield the outer electrons from the
                                                              nuclear charge in each element increases,
attractive force of the positive nucleus.
                                                              this increases the attractive force acting on
as the radius of the atom increases, the distance
                                                              the outer most electrons.
between the nucleus and the outer electron
                                                              there is no significant increase in shielding as
increases and therefore the force of attraction
                                                              each successive electron enters the same
between the nucleus and outer most electrons is
                                                              energy level as the one before.
                                                              Overall more energy is needed to remove the
These factors mean that less energy is needed
                                                              first electron from its outermost shell.
to remove the first electron from an atom at the
bottom of the group compared to one at the top
of the group.
        Amount of Substance
•   Amount of Substance
•   Calculations
•   Balancing Equations
•   Reacting Quantities
        Amount of Substance
• Different atoms have different masses. 1g of
  carbon has far fewer atoms than 1g of hydrogen
  atoms. Chemists need a method of quantifying
• We use a quantity called the amount of
  substance which is measured in moles.
• One mole contains 6.02 x 1023 particles.
• The relative atomic mass is the mass of one
  mole of that element’s atoms.
• The relative molecular mass is the mass of one
  mole of molecules
• Amount of substance, n           • The Ideal Gas equation
           – n = mass
                  Mr               • Providing that the
                                     pressure and temperature
• Solution calculations              of gases are the same,
   – The concentration of            equal volumes of two
     solutions are measured in       gases can be assumed to
     mol dm3
       • c= concentration in dm3     have the same number of
       • V = volume in cm3           moles.
              – n = Vc                – pV = nRT
                  1000                   • p is the pressure in Pa
                                         • V is the volume in m3
                                         • T is the temperature in
 P                                         Kelvin
 r                                       • R= 8.31 JK-1 mol-1
         Balancing Equations
• Chemical reactions involve the rearrangement of
  atoms not the making or destroying of atoms.
• It is necessary to make sure that you have the
  same amount of atoms on both sides of the
• State symbols can also be added to show the
  physical condition of the reactants and products
• (s) – solid, (l) – liquid, (g) – gas, (aq) – aqueous
             Reacting Quantities
• The numbers in a balanced equation give the ratio of the amount of
  each substance in the reaction. We can use this information to
  calculate quantities of reactants or products.
• 50g of CaCO3 are heated how much CaO will be formed
• First write a balanced equation:
    – CaCO3(s)  CaO(s) + CO2(g)
• Then calculate the Mr of the substances we are interested in:
    – CaCO3 40 + 12 + (3 x 16) = 100
    – CaO 40 + 16 =56
• Calculate the number of moles of CaCO3 used.
    – n=mass/Mr 50/100 = 0.5 mol
• From the equation we can see that one mole of CaCO3 produces
  one mole of CaO. Therefore 0.5 mol of CaCO3 produces 0.5 mol of
• Finally calculate the mass of 0.5 mol of CaO:
    – n=mass/Mr, mass=Mr x n, 0.5 x 56 = 28g
• Therefore 50g of CaCO3 produces 28g CaO
•   The nature of bonds
•   Bond polarity and the polarisation of ions
•   Intermolecular forces
•   Hydrogen Bonding
                              The Nature of Bonds
    Covalent Bonding

                      H                           •    When non-metals react together both atoms need to gain
                                                       electrons to obtain a full shell of electrons they do this by
                                                       forming a covalent bond.
         H            C           H               •    The atoms are held together by shared pairs of electrons from
                                                       the highest energy level of both atoms.
    Ionic bonding
                                                                  Atoms lose or gain electrons to attain a complete outer
                          +                             _         shell of electrons.
                                                                  An ionic bond is formed when electrons are lost and
                                                                  gained by two or more atoms.
         Na                            Cl                         When atoms lose electrons they become positive ions,
                                                                  when they gain electrons they become negative ions.
                                                                  It is the electrostatic forces of attraction which hold the
                                                                  ions together
Metallic Bonding
                                                                 In metals, positive metal ions are held
                              -              -
-   +-        -       +               -+              +     - - ++        - +      -    +
                                                                 together by electron clouds. These
      +           -       +              +    -         +-       electrons-are+ to-move through the
                                                                               free        +
     -    +                   +       - -+            - +                        +-          +
                                                                    + this is why metals conduct
              -       -           +          +             + - structure, - +
                                                                      +                    - +
               Bond Polarity and the
                Polarisation of Ions
• In reality not all bonds are perfectly covalent or ionic. To explain why
  we have to define a concept called electronegativity.
• Electronegativity is the ability of an atom to attract the bonding
• In hydrogen fluoride the fluorine atoms are much more
  electronegative than the hydrogen. It pulls the electrons toward it
  creating what is called a polar bond.
• Ionic bonds can also show polarity, this can happen if the electron
  cloud is distorted by strong charges on one of the ions. If a cation is
  highly charged it will exert a strong electrostatic attraction on the
  anion and distort the electron cloud. If the anion has a large electron
  cloud it will be easily distorted. If the electron cloud is distorted there
  will be electron density between the two ions giving the bond some
  covalent characteristics
• Molecules with asymmetric charge distribution are said to be polar
  molecules and to possess a dipole
                Intermolecular Forces
•   Permanent dipole-
    permanent dipole
    interactions occur                      d+ H         Cl d-
    between polar molecules.
    This happens when the
    negative end of one                                               Weak electrostatic forces
    molecule is attracted to
    the positive end of
    another. This force is                  d- Cl         H d+
    much weaker than
    intramolecular bonding

    •Temporary dipole –induced dipole interactions exist between non-polar
    molecules and monatomic species such as the noble gases. The
    distribution of the electron cloud on a molecule is not constant and at any
    given time it can asymmetric. this confers a temporary asymmetry on the
    charge distribution. The molecule is said to possess a temporary dipole,
    this temporary dipole can induce another temporary dipole in an adjacent
    molecule. There is a resulting weak electrostatic force between the two
                         Hydrogen Bonding
                          • Hydrogen bonds are a special case of
                            permanent dipole-permanent dipole
         d-                 bonding.
      O                   • It exists where an electronegative
                            element such as oxygen, chlorine
              d+            fluorine or nitrogen is bonded to
               d-           hydrogen.
              O           • Hydrogen bonding causes stronger
                            intermolecular bonds than would
                    d+      otherwise be predicted this increases
                            the boiling point of substances such as
     O                      water.
• Chemists classify elements according to their position in the periodic
• Periodicity is the term used to describe the repeating pattern of
  properties observed within the periodic table.

                                              H                                     He

          Li   Be                                            B    C    N   O   F    Ne

          Na   Mg                                            Al   Si   P   S   Cl   Ar

          K    Ca Sc Ti      V    Cr Mn Fe Co Ni       Cu Zn Ga Ge As Se Br         Kr

          Rb   Sr   Y   Zr   Nb Mo Tc      Ru Rh Pd Ag Cd In      Sn Sb Te I        Xe

          Cs   Ba La Hf      Ta W    Re Os Ir     Pt   Au Hg Tl   Pb Bi    Po At    Rn

          Fr   Ra Ac

       s- block                  d-block                          p-block
                 Trends in Group 2 Compounds
• Progressing down group 2 the atomic radius increases due to the
  extra shell of electrons for each element.
• Going down the group the first ionisation energy decreases there is
  more shielding between the nucleus and the outer electrons and the
  distance between the nucleus and the outer electron increases and
  therefore the force of attraction between the nucleus and outer most
  electrons is reduced.
• Generally the melting point of the metals decreases down the group
  this is because as the metal ions get larger the distance between the
  bonding electrons and the positive nucleus gets larger and reduces
  the overall attraction between the two. For similar reasons the
  electronegativity decreases.
• The reactions of the elements with water become more vigorous
  down the group. When they do react they produce hydroxides and
• The solubilities of the hydroxides of the elements increase going
  down the group.
• The solubilities of the sulphates of the elements decreases down the
• Barium sulphate is insoluble and is used as a qualitative test to
  identify sulphate ions.
 Trends in Period Three of the Periodic Table
Property            Trend from        Explanation
                    left to right
Atomic radius       decreases         because the nuclear charge
First ionisation    increases         because the nuclear charge
energy                                increases
Electronegativity   increases         because the nuclear charge
Electrical          Increases until   because the metals have an
conductivity        the non metals    increased number of delocalised
Boiling point and   Increases until   Because these properties depend
                    the middle then   on the forces between the particles.
Melting point
                    decreases         This depends on the structure of the
                                      element which varies from metallic
                                      to giant covalent to simple
• Atomic Structure
   – We consider atoms to be formed from three fundamental
     particles, we can determine the relative atomic mass using a
     mass spectrometer.
• Amount of Substance
   – Chemists use this concept to count atoms. Using this concept
     we can calculate reacting quantities, in a given reaction.
• Bonding
   – Bonding within molecules can be described as covalent, ionic or
     metallic. Often a bond is a hybrid between ionic and covalent.
• Periodicity
   – Trends within the periodic table can frequently be explained by
     concepts of nuclear charge and shielding.