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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
Atomic
Particle
Electron
+
+
Relative
Charge
neutron
1
0
proton
1
+1
negligible
-1
electron
Proton
Relative
Mass
Isotopes
• Atoms of the same
element always have the
same number of protons.
• Isotopes are atoms of the
same element but with
different number of
neutrons. This gives rise
to different mass
numbers.
• Relative abundance is the
amount of each isotope as
the percentage for that
element occurring on the
Earth
Mass number: this is the number of
protons and neutrons (A)
4
2
He
Helium
Atomic number: this is the
number of protons (Z)
Mass Spectrometer
The
sample is
put in via
an injector
An
electron
gun
ionises
the atoms
Ionisation
Charged
plates
accelerate
positive
ions
acceleration
•The mass spectrum gives the relative
masses of each isotope and the
abundance of each isotope.
•The relative atomic mass is the
weighted average mass of an atom of an
element compared with 1/12 of the mass
of 12C.
•We can calculate this from mass
spectra
A curved magnet
deflects the
positive ions. The
lighter ions are
deflected more
deflection
The ions are detected
to produce a mass
spectrum
detection
Example: Copper
69% of copper atoms have a
relative mass of 63.
31% of copper atoms have a
relative mass of 65.
The weighted average is calculated
as follows:
(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
configuration
H
1s
He
Li
Be
B
C
1s2
1s22s
1s22s2
1s22s22p1
1s22s22p2
N
1s22s22p3
O
1s22s22p4
F
1s22s22p5
Element
Electron configuration
Ne
1s22s22p6
Na
1s22s22p63s1
Mg
1s22s22p63s2
Al
1s22s22p63s23p3
Si
1s22s22p63s23p3
P
1s22s22p63s23p3
S
1s22s22p63s23p4
Cl
1s22s22p63s23p5
Ar
1s22s22p63s23p6
Trends in Ionisation Energies
1000
First ionisation energies of period 3
elements
Energy (Kj/mol)
Energy (Kj/mol)
The first ionisation energies of group 2
electrons
800
600
400
200
0
2000
1500
1000
500
0
Be
Mg
Ca
Sr
Ba
Group 2 elements
Going down a group in the periodic table
there are more filled energy levels between the
nucleus and the outer most electrons
these shield the outer electrons from the
attractive force of the positive nucleus.
as the radius of the atom increases, 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.
These factors mean that less energy is needed
to remove the first electron from an atom at the
bottom of the group compared to one at the top
of the group.
Na
Mg
Al
Si
P
S
Cl
Ar
Period 3 elements
Going across a period of elements:
there are more protons in each nucleus so the
nuclear charge in each element increases,
this increases the attractive force acting on
the outer most electrons.
there is no significant increase in shielding as
each successive electron enters the same
energy level as the one before.
Overall more energy is needed to remove the
first electron from its outermost shell.
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
– n = mass
Mr
• Solution calculations
– The concentration of
solutions are measured in
mol dm3
• c= concentration in dm3
• V = volume in cm3
– n = Vc
1000
P
r
• The Ideal Gas equation
• Providing that the
pressure and temperature
of gases are the same,
equal volumes of two
gases can be assumed to
have the same number of
moles.
– pV = nRT
• p is the pressure in Pa
• V is the volume in m3
• T is the temperature in
Kelvin
• 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
H
•
H
C
•
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.
The atoms are held together by shared pairs of electrons from
the highest energy level of both atoms.
H
Ionic bonding
_
+
Cl
Na
Metallic 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.
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
In metals, positive metal ions are held
- clouds.
+
+ These
- +together- by+electron
++through the
+
electrons
-are+free to-move
- +
+
+
- metals+conduct
+ - structure,
+ this- is why
+
- +
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 dipolepermanent dipole
interactions occur
between polar molecules.
This happens when the
negative end of one
molecule is attracted to
the positive end of
another. This force is
much weaker than
intramolecular bonding
d+ H
Cl dWeak electrostatic forces
d- Cl
H d+
•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
Hydrogen Bonding
d-
O
H
d+
H
d+
d-
O
H
d+
H
d+
d-
O
H
d+
d-
O
H
d+
H
d+
H
d+
• Hydrogen bonds are a special case of
permanent dipole-permanent dipole
bonding.
• It exists where an electronegative
element such as oxygen, chlorine
fluorine or nitrogen is bonded to
hydrogen.
• Hydrogen bonding causes stronger
intermolecular bonds than would
otherwise be predicted this increases
the boiling point of substances such as
water.
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
Cu Zn Ga Ge As Se Br
Kr
Rb
Sr
Nb Mo Tc
Cs
Ba La Hf
Fr
Ra Ac
s- block
Y
Zr
Cr Mn Fe Co Ni
Ta W
Ru Rh Pd Ag Cd In
Re Os Ir
d-block
Pt
Au Hg Tl
Sn Sb Te I
Xe
Pb Bi
Rn
Po At
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
left to right
Explanation
Atomic radius
decreases
because the nuclear charge
increases
First ionisation
energy
increases
because the nuclear charge
increases
Electronegativity
increases
because the nuclear charge
increases
Electrical
conductivity
Increases until
the non metals
because the metals have an
increased number of delocalised
electrons
Boiling point and
Melting point
Increases until
the middle then
decreases
Because these properties depend
on the forces between the particles.
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.