IB Chemistry (unit 2) ATOMIC THEORY

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IB Chemistry (unit 2)ATOMIC THEORY
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Atomic Structure - Recap

• Questions
• Define the following words
• Atom
• Element
• Molecule
• Compound

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Atomic Structure

• Atoms are very small 10-10 metres
• All atoms are made up of three sub-atomic
  particles protons, neutrons and electrons

• The protons and neutrons form a small positively
  charged nucleus
• The electrons are in energy levels outside the
  nucleus

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Atomic Structure

• The actual values of the masses and charges of
  the sub-atomic particles are shown below

• A meaningful way to consider the masses of the
  sub-atomic particles is to use relative masses

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Atomic Structure - Definitions

• Atomic number (Z) is the number of protons in the
  nucleus of an atom. It is also known as the
  proton number.
• N.B. No. of protons always equals the no. of
  electrons in any neutral atom of an element.
• Mass number (A) is the sum of the number of
  protons and the number of neutrons in the nucleus
  of an atom.

• So how can you work out the number of neutrons in
  an atom?

No. of neutrons Mass number atomic number
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Atomic Structure - Example

• So how can you work out the number of neutrons in
  an atom?
• Example

No. of neutrons Mass number atomic number
No. of neutron Mass No. Atomic No.
23 11 12
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Atomic Structure - Questions

• What are the three sub atomic particles that make
  up the atom?
• Draw a representation of the atom and labelling
  the sub-atomic particles.
• Draw a table to show the relative masses and
  charges of the sub-atomic particles.
• State the atomic number, mass number and number
  of neutrons of a) carbon, b) oxygen and c)
  selenium.
• Which neutral element contains 11 electrons and
  12 neutrons?

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Atomic Structure - Questions

• 5. Copy and complete the following table

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Summary Slide

• All atomic masses are relative to the mass of
  carbon-12.
• Eg one hydrogen atom weighs 1/12 the mass of a
  carbon-12 atom.

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Isotopes

• Isotopes are atoms of the same element with the
  same atomic number, but different mass numbers,
  i.e. they have different numbers of neutrons.

Each atom of chlorine contains the following
17 protons 17 electrons 18 neutrons
17 protons 17 electrons 20 neutrons
The isotopes of chlorine are often referred to as
chlorine-35 and chlorine-37
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Isotopes

• Isotopes of an element have the same chemical
  properties because they have the same number of
  electrons. When a chemical reaction takes place,
  it is the electrons that are involved in the
  reactions.
• However isotopes of an element have the slightly
  different physical properties because they have
  different numbers of neutrons, hence different
  masses.
• The isotopes of an element with fewer neutrons
  will have
• Lower masses faster rate of diffusion
• Lower densities lower melting and boiling
  points

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Isotopes - Questions

• Explain what isotopes using hydrogen as an
  example.
• One isotope of the element chlorine, contains 20
  neutrons. Which other element also contains 20
  neutrons?
• State the number of protons, electrons and
  neutrons in
• a) one atom of carbon-12
• b) one atom of carbon-14
• c) one atom of uranium-235
• d) one atom of uranium-238

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Mass Spectrometer

• The mass spectrometer is an instrument used
• To measure the relative masses of isotopes
• To find the relative abundance of the isotopes in
  a sample of an element

When charged particles pass through a magnetic
field, the particles are deflected by the
magnetic field, and the amount of deflection
depends upon the mass/charge ratio of the charged
particle.
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Mass Spectrometer 5 Stages

• Once the sample of an element has been placed in
  the mass spectrometer, it undergoes five stages.
• Vaporisation the sample has to be in gaseous
  form. If the sample is a solid or liquid, a
  heater is used to vaporise some of the sample.

X (s) ? X (g) or X (l) ? X (g)
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Mass Spectrometer 5 Stages

• Ionisation sample is bombarded by a stream of
  high-energy electrons from an electron gun, which
  knock an electron from an atom. This produces
  a positive ion

X (g) ? X (g) e-

• Acceleration an electric field is used to
  accelerate
• the positive ions towards the magnetic field.
  The
• accelerated ions are focused and passed
  through a
• slit this produces a narrow beam of ions.

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Mass Spectrometer 5 Stages

• Deflection

The accelerated ions are deflected into the
magnetic field. The amount of deflection is
greater when the mass of the positive ion is
less the charge on the positive ion is
greater the velocity of the positive ion is
less the strength of the magnetic field is
greater
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Mass Spectrometer

• If all the ions are travelling at the same
  velocity and carry the same charge, the amount of
  deflection in a given magnetic field depends upon
  the mass of the ion.
• For a given magnetic field, only ions with a
  particular relative mass (m) to charge (z) ration
  the m/z value are deflected sufficiently to
  reach the detector.

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Mass Spectrometer

• Detection ions that reach the detector cause
  electrons to be released in an ion-current
  detector
• The number of electrons released, hence the
  current produced is proportional to the number of
  ions striking the detector.
• The detector is linked to an amplifier and then
  to a recorder this converts the current into a
  peak which is shown in the mass spectrum.

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Atomic Structure Mass Spec

• Name the five stages which the sample undergoes
  in the mass spectrometer and make brief notes of
  what you remember under each stage.
• Complete Exercise 4, 5 and 6 in the handbook.
  Any incomplete work to be completed and handed in
  for next session.

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Atomic Structure Mass Spec

• Isotopes of boron

m/z value 11 10
Relative abundance 18.7 81.3
Ar of boron (11 x 18.7) (10 x 81.3)
(18.7 81.3) 205.7 813
100 1018.7 10.2
100
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Mass Spectrometer Questions

• A mass spec chart for a sample of neon shows that
  it contains
• 90.9 20Ne
• 0.17 21Ne
• 8.93 22Ne
• Calculate the relative atomic mass of neon
• You must show all your working!

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Mass Spectrometer Questions

• 90.9 20Ne
• 0.17 21Ne
• 8.93 22Ne

Ar 20.18
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Mass Spectrometer Questions

• Calculate the relative atomic mass of lead
• You must show all your working!

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Mass Spectrometer Questions

• 1.5 204Pb
• 23.6 206Pb
• 22.6 207Pb
• 52.3 208Pb

Ar 207.24
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Energy Levels

• Electrons go in shells or energy levels. The
  energy levels are called principle energy levels,
  1 to 4.
• The energy levels contain sub-levels.

Principle energy level Number of sub-levels
1 1
2 2
3 3
4 4
These sub-levels are assigned the letters, s, p,
d, f
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Energy Levels

• Each type of sub-level can hold a different
  maximum number of electron.

Sub-level Maximum number of electrons
s 2
p 6
d 10
f 14
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Energy Levels

• The energy of the sub-levels increases from s to
  p to d to f. The electrons fill up the lower
  energy sub-levels first.

Looking at this table can you work out in what
order the electrons fill the sub-levels?
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Energy Levels

• Lets take a look at the Periodic Table to see
  how this fits in.

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Electronic Structure

• So how do you write it?

1s2
Example For magnesium 1s2, 2s2, 2p6, 3s2
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Electronic Structure

• The electronic structure follows a pattern the
  order of filling the sub-levels is 1s, 2s, 2p,
  3s, 3p
• After this there is a break in the pattern, as
  that the 4s fills before 3d.
• Taking a look at the table below can you work
  out why this is?

• This is because the 4s
• sub-level is of
• lower energy than the
• 3d sub-level.

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Electronic Structure

• The order in this the energy levels are filled is
  called the Aufbau Principle.
• Example (Sodium 2, 8, 1)

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Electronic Structure

• There are two exceptions to the Aufbau principle.
• The electronic structures of chromium and copper
  do not follow the pattern they are anomalous.
• Chromium 1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s1
• Copper 1s2, 2s2, 2p6, 3s2. 3p6, 3d10, 4s1

• Write the electronic configuration for the
  following elements
• hydrogen c) oxygen e) copper
• carbon d) aluminium f) fluorine

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Electronic Structure of ions

• When an atom loses or gains electrons to form an
  ion, the electronic structure changes
• Positive ions formed by the loss of e-

1s2 2s2 2p6 3s1 ?
1s2 2s2 2p6
Na atom
Na ion

• Negative ions formed by the gain of e-

1s2 2s2 2p4 ?
1s2 2s2 2p5
O atom
O- ion
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Electronic Structure of transition metals

• With the transition metals it is the 4s
  electrons that are lost first when they form
  ions
• Titanium (Ti) - loss of 2 e-

1s2 2s2 2p6 3s2 3p6 3d2 4s2 ?
1s2 2s2 2p6 3s2 3p6 3d2
Ti atom
Ti2 ion

• Chromium (Cr) - loss of 3 e-

1s2 2s2 2p6 3s2 3p6 3d3
1s2 2s2 2p6 3s2 3p6 3d5 4s1 ?
Cr atom
Cr3 ion
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Electronic Structure - Questions

• Give the full electronic structure of the
  following positve ions
• a) Mg2 b) Ca2 c) Al3
• Give the full electronic structure of the
  negative ions
• a) Cl- b) Br- c) P3-

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Electronic Structure - Questions

• Copy and complete the following table

Atomic no. Mass no. No. of protons No. of neutrons No. of electrons Electronic structure
Mg 12 1s2 2s2 2p6 3s2
Al3 27 10
S2- 16 16
Sc3 21 45
Ni2 30 26
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Orbitals

• The energy sub levels are made up of orbitals,
  each which can hold a maximum of 2 electrons.
• Different sub-levels have different number of
  orbitals

Sub-level No. of orbitals Max. no. of electrons
s 1 2
p 3 6
d 5 10
f 7 14
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Orbitals

• The orbitals in different sub-levels have
  different shapes

• s orbitals

• p orbitals

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Orbitals

• Within a sub-level, the electrons occupy orbitals
  as unpaired electrons rather than paired
  electrons. (This is known as Hunds Rule).
• We use boxes to represent orbitals

?
?
?
?
Electronic structure of carbon, 1s2, 2s2, 2p2
?
?
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Orbitals

• The arrows represent the electrons in the
  orbitals.
• The direction of arrows indiactes the spin of the
  electron.
• Paired electrons will have opposite spin, as this
  reduces the mutual repulsion between the paired
  electrons.

Electronic structure of carbon, 1s2, 2s2, 2p2
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Orbitals

• Using boxes to represent orbitals, give the full
  electronic structure of the following atoms
• a) lithium b) fluorine c) potassium
• d) nitrogen e) oxygen

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Orbitals

• Using boxes to represent orbitals, give the full
  electronic structure of the following atoms
• a) lithium b) fluorine c) potassium
• d) nitrogen e) oxygen

Electronic structure of lithium 1s2, 2s1
?
?
?
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Orbitals

• Using boxes to represent orbitals, give the full
  electronic structure of the following atoms
• a) lithium b) fluorine c) potassium
• d) nitrogen e) oxygen

Electronic structure of fluorine 1s2, 2s2, 2p5
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Orbitals

• Using boxes to represent orbitals, give the full
  electronic structure of the following atoms
• a) lithium b) fluorine c) potassium
• d) nitrogen e) oxygen

Electronic structure of potassium 1s2, 2s2, 2p6,
3s2, 3p6, 4s1
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Orbitals

• Using boxes to represent orbitals, give the full
  electronic structure of the following atoms
• a) lithium b) fluorine c) potassium
• d) nitrogen e) oxygen

Electronic structure of nitrogen 1s2, 2s2, 2p3
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Orbitals

• Using boxes to represent orbitals, give the full
  electronic structure of the following atoms
• a) lithium b) fluorine c) potassium
• d) nitrogen e) oxygen

Electronic structure of oxygen 1s2, 2s2, 2p4
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Ionisation Energy

• Ionisation of an atom involves the loss of an
  electron to form a positive ion.
• The first ionisation energy is defined as the
  energy required to remove one mole of electrons
  from one mole of atoms of a gaseous element.
• The first ionisation energy of an atom can be
  represented by the following general equation
• X(g) ? X e- ?H ve
• Since all ionisations requires energy, they are
  endothermic processes and have a positive
  enthalpy change (?H) value.

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Ionisation Energy

• The value of the first ionisation energy depends
  upon two main factors
• The size of the nuclear charge
• The energy of the electron that has been removed
  (this depends upon its distance from the nucleus)

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Ionisation Energy

• As the size of the nuclear charge increases the
  force of the attraction between the negatively
  charged electrons and the positively charged
  nucleus increases.

Small nuclear charge ?
Large nuclear charge ?
?
?
Small force of attraction ?
Large force of attraction ?
Smaller ionisation energy
Greater ionisation energy
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Ionisation energy

• As the energy of the electron increases, the
  electron is farther away from the nucleus. As a
  result the force of attraction between the
  nucleus and the electron decreases.

Electrons further away from positive nucleus ?
Electrons closer to positive nucleus ?
Large force of attraction ?
Small force of attraction ?
Greater ionisation energy
Smaller ionisation energy
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Ionisation energy - Questions

• Write an equation to represent the first
  ionisation of
• a) aluminium
• b) lithium
• c) sodium

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Trends across a Period

• Going across a period, the size of the 1st
  ionisation energy shows a general increase.
• This is because the electron comes from the same
  energy level, but the size of the nuclear charge
  increases.

Going across a Period
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Trends across a Period (2 exceptions)

• The first ionisation of Al is less than that of
  Mg, despite the increase in the nuclear charge.
• The reason for this is that the outer electron
  removed from Al is in a higher sub-level the
  electron removed from Al is a 3p electron,
  whereas that removed from Mg is a 3s.

Electronic structure Ionisation energy/kJ mol-1
Na 1s2, 2s2, 2p6, 3s1 494
Mg 1s2, 2s2, 2p6, 3s2 736
Al 1s2, 2s2, 2p6, 3s2, 3p1 577
Si 1s2, 2s2, 2p6, 3s2, 3p2 786
P 1s2, 2s2, 2p6, 3s2, 3p3 1060
S 1s2, 2s2, 2p6, 3s2, 3p4 1000
Cl 1s2, 2s2, 2p6, 3s2, 3p5 1260
Ar 1s2, 2s2, 2p6, 3s2, 3p6 1520
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Trends across a Period (2 exceptions)

• The first ionisation energy of S is less than
  that of P, despite the increase in the nuclear
  charge.
• In both cases the electron removed is from the 3p
  sub-level. However the 3p electron removed from
  S is a paired electron, whereas the 3p electron
  removed from P is an unpaired electron.
• When the electrons are paired the extra mutual
  repulsion results in less energy being required
  to remove an electron, hence a reduction in the
  ionisation energy.

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Trends across a Period - Questions

• There is a break in this general trend going
  across a Period.
• Look at the table below and point out where the
  break in the the trend is and try to give an
  explanation.

Electronic structure Ionisation energy/kJ mol-1
Na 1s2, 2s2, 2p6, 3s1 494
Mg 1s2, 2s2, 2p6, 3s2 736
Al 1s2, 2s2, 2p6, 3s2, 3p1 577
Si 1s2, 2s2, 2p6, 3s2, 3p2 786
P 1s2, 2s2, 2p6, 3s2, 3p3 1060
S 1s2, 2s2, 2p6, 3s2, 3p4 1000
Cl 1s2, 2s2, 2p6, 3s2, 3p5 1260
Ar 1s2, 2s2, 2p6, 3s2, 3p6 1520
Clue which sub-level (s, p, d or f is the outer
electron in?
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Trends across a Period - Questions

• Now take a look at the graph below

1. Explain what the graph shows in as much detail as
   possible
2. There is one other break in the general pattern
   going across a Period. What is it and explain
   why that is.

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Trends down a Group

• Ionisation energy decreases going down a Group.
• Going down a Group in the Periodic Table, the
  electron removed during the first ionisation is
  from a higher energy level and hence it is
  further from the nucleus.
• The nuclear charge also increases, but the effect
  of the increased nuclear charge is reduced by the
  inner electrons which shield the outer electrons.

Down the Group
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Ionisation energy - Questions

1. Explain why sodium has a higher first ionisation
   energy than potassium.
2. Explain why the first ionisation energy of boron
   is less than that of beryllium.
3. Why does helium have the highest first ionisation
   energy of all the elements?
4. Complete Tasks

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Successive Ionisation energy

• Definition 2nd i.e.
• The energy per mole for the process
• X(g) X2(g) e-
• And so on for further successive ionisation
  energies

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Successive Ionisation energy

• Successive i.es increases because electrons are
  being removed from increasingly positive ions.
• Therefore, nuclear attraction is greater.
• Large jumps seen when electron is removed form a
  new sublevel closer to the nucleus

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Successive Ionisation energy
Large increase between 4th and 3rd shells
electron closer to nucleus
2nd i.e higher than first electron has greater
pull from nucleus
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Electron Affinity

• Energy Change per mole for
• X (g) e- X-(g)
• That is, for the gaseous atoms to gain an
  electron to form anions

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Electron Affinity
The first e.a is negative (exothermic) because
the electron is attracted to the positive charge
on the atoms nucleus. The second e.a is
positive (endothermic) because an electron is
being added to an ion which is already negative
repulsion occurs