# Atomic Structure_ Bonding and Periodicity

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```					Atomic Structure, Bonding and
Periodicity
Contents
•   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.

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

Proton
Isotopes
• 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
numbers.
2   He
Helium
• 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
Earth
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
spectra
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
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
600
1000
400
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.
reduced.
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
atoms.
• 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
Calculations
• 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
equation.
• 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
CaO.
• 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
Bonding
•   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.
H
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, - +
+                    - +
electricity.
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
electrons.
• 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
molecules
d-
Hydrogen Bonding
O
• Hydrogen bonds are a special case of
H
d+
H
d+
permanent dipole-permanent dipole
d-                 bonding.
O                   • It exists where an electronegative
element such as oxygen, chlorine
H
d+
H
d+            fluorine or nitrogen is bonded to
d-           hydrogen.
O           • Hydrogen bonding causes stronger
intermolecular bonds than would
H
d+
H
d+      otherwise be predicted this increases
d-
the boiling point of substances such as
O                      water.
H
d+
H
d+
Periodicity
• Chemists classify elements according to their position in the periodic
table.
• 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
hydrogen.
• The solubilities of the hydroxides of the elements increase going
down the group.
• The solubilities of the sulphates of the elements decreases down the
group.
• 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
increases
First ionisation    increases         because the nuclear charge
energy                                increases
Electronegativity   increases         because the nuclear charge
increases
Electrical          Increases until   because the metals have an
conductivity        the non metals    increased number of delocalised
electrons
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
molecular.
Summary
• 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.

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