unit_1_mod_1_atomic_structure_and_periodic_table.doc

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ATOMIC STRUCTURE AND THE PERIODIC TABLE

Theoretical change with respect to Daltons atomic theory

1. In 1803, atomic theory was revived by John Daltona) matter is made up of tiny particles called atoms which cannot be created,

destroyed or split

b) all atoms of one element are identical:- same mass and same chemical

properties

c) a chemical reaction consists of rearranging atoms from one combination to

another.

d) When elements combine to form compounds, small whole numbers of atoms

form molecules.

However this was proved to be not entirely correct. Atoms have been spl

it as

well as created i.e. nuclear reactions. Also there are isotopes, meaning that not

all atoms of an element are identical.

Therefore theory was forced to CHANGE in regards to these observations

contradicting to the theory

put forward by Dalton.

The distribution of charge and mass in an atom

Particle Location Mass Charge

Electron Orbitals 1/1837 unit -1 unit

Proton Nucleus 1 unit +1 unit

Neutron Nucleus 1 unit 0

Atomic Structure page 1 of 20

A unit is one atomic mass unit = 1.67 x 10-27


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Terminology

Term

Definition

Atomic/proton number Number of protons in a nucleus of an atomNucleon/mass number Sum of the number of protons and neutrons in the

nucleus of an atom

Nuclide Any atomic species of which the proton number and

nucleon number are specified e.g. 126C and94B are

nuclides

Isotopes Nuclides of the same element or atoms of the same

element with different mass numbers

NB isotopes have the same chemical properties but

different physical propertiesRelative atomic mass Mass of an atom based on a scale such that the C-12

isotope has a mass of 12.00 units

relative atomic mass

= mass of 1 atom of an element x 12

mass of 1 atom of carbon-12



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Phenomenon of radioactivity

Radiation is the spontaneous decay of unstable atoms with the emission of either

alpha, beta or gamma radiation.

Alpha decay is a type ofradioactive decay in which anatomic nucleusemits an

alpha particle (two protons and two neutrons bound together into a particle

identical to a helium nucleus) and transforms (or 'decays') into an atom with a

mass number4 less andatomic number2 less.

For example:

although this is typically written as:

Beta decay is a type ofradioactive decay in which a fast moving ele

ctrons is

emitted. The new atom has no change in mass number but an atomic number

increases by 1.

Gamma rays orgamma-ray (denoted as ) are forms ofelectromagneticradiation(EMR) orlight emissions of a specific frequency produced from sub-

atomicparticle interaction, such aselectron-positron annihilationand

radioactive decay. There is no change in atomic or mass number of the atom.

Band of stability (n/p ratio)


Most elements have isotopes. For stableisotopes, an interesting plot arises when

the number of neutrons is plotted versus

the number of protons.

Because the plot shows only the stable

isotopes, this graph is often called theNuclear Belt of Stability.

The plot indicates that lighter nuclides

(isotopes) are most stable when theneutron/proton ratio is 1/1. This is the

case with any nucleus that has up to 20protons.

As theatomic numberincreases beyond

20, a different trend becomesapparent. In this range, it appears that a

stable nucleusis able to accommodatemore neutrons. Stable isotopes have a

higherneutrontoprotonratio, rising to1.5/1 for elements having atomic

numbers between 20 and 83.
http://en.wikipedia.org/wiki/Radioactivityhttp://en.wikipedia.org/wiki/Radioactivityhttp://en.wikipedia.org/wiki/Atomic_nucleushttp://en.wikipedia.org/wiki/Atomic_nucleushttp://en.wikipedia.org/wiki/Atomic_nucleushttp://en.wikipedia.org/wiki/Alpha_particlehttp://en.wikipedia.org/wiki/Atomic_weighthttp://en.wikipedia.org/wiki/Atomic_numberhttp://en.wikipedia.org/wiki/Atomic_numberhttp://en.wikipedia.org/wiki/Radioactivehttp://en.wikipedia.org/wiki/Radioactivehttp://en.wikipedia.org/wiki/Gammahttp://en.wikipedia.org/wiki/Electromagnetic_radiationhttp://en.wikipedia.org/wiki/Electromagnetic_radiationhttp://en.wikipedia.org/wiki/Electromagnetic_radiationhttp://en.wikipedia.org/wiki/Electromagnetic_radiationhttp://en.wikipedia.org/wiki/Photonhttp://en.wikipedia.org/wiki/Atomhttp://en.wikipedia.org/wiki/Atomhttp://en.wikipedia.org/wiki/Atomhttp://en.wikipedia.org/wiki/Electron-positron_annihilationhttp://en.wikipedia.org/wiki/Electron-positron_annihilationhttp://en.wikipedia.org/wiki/Electron-positron_annihilationhttp://en.wikipedia.org/wiki/Radioactive_decayhttp://en.wikipedia.org/wiki/Radioactive_decayhttp://def%28%27/Glossary/glossaryterm.aspx?word=Graph%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Ratio%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Nucleus%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Atomic%20Number%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Atomic%20Number%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Atomic%20Number%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Nucleus%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Nucleus%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Neutron%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Neutron%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Neutron%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Proton%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Proton%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Proton%27,%20500,%20500);http://en.wikipedia.org/wiki/Atomic_nucleushttp://en.wikipedia.org/wiki/Alpha_particlehttp://en.wikipedia.org/wiki/Atomic_weighthttp://en.wikipedia.org/wiki/Atomic_numberhttp://en.wikipedia.org/wiki/Radioactivehttp://en.wikipedia.org/wiki/Gammahttp://en.wikipedia.org/wiki/Electromagnetic_radiationhttp://en.wikipedia.org/wiki/Electromagnetic_radiationhttp://en.wikipedia.org/wiki/Photonhttp://en.wikipedia.org/wiki/Atomhttp://en.wikipedia.org/wiki/Atomhttp://en.wikipedia.org/wiki/Electron-positron_annihilationhttp://en.wikipedia.org/wiki/Radioactive_decayhttp://def%28%27/Glossary/glossaryterm.aspx?word=Graph%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Ratio%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Nucleus%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Atomic%20Number%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Nucleus%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Neutron%27,%20500,%20500);http://def%28%27/Glossary/glossaryterm.aspx?word=Proton%27,%20500,%20500);http://en.wikipedia.org/wiki/Radioactivity


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Uses of radioisotopes

1. radiocarbon dating 2. smoke detectors

3. pacemakers 4. medical uses i.e. trac

ers or chemotherapy

5. irradiation in pest control

Calculations of relative atomic mass from isotopic data

Arof an element = sum of (abundances x mass number of all of the isotopes of

an element)

e.g.zirconium-90 51.5% zirconium-91 11.2% zirconium-92 17.1%

zirconium-94 17.4 % zirconium-96 2.8%

Arzirconium = (51.5 x 90) + (11.2 x 91) + (17.1 x 92) + (17.4 x 94) + (2.8 x 96)

= 9131.8

The average mass of these 100 atoms would be 9131.8 / 100 = 91.3 (to 3

significant figures).

91.3 is the relative atomic mass of zirconium.

Evidence of discrete energy levels using emission spectra



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An element's emission spectrum is the relative intensity ofelectromagnetic

radiation of each frequency it emits when it is heated (or more generally when it

is excited).

When the electronsin the elementare excited, they jump to higher energylevels. As the electrons leave the excited state and fall back down, energy is

emitted, the wavelength of which refers to the discrete lines of the emission

spectrum. If discrete energy levels were NOT present, no lines could EVER be

formed in an emission spectrum.

e.g. emission spectrum of hydrogen

Emission spectrum of hydrogen

When a gaseous hydrogen atom i

n its ground state is excited by an input of

energy, its electron is 'promoted' from the lowest energy level to one of higher

energy. The atom does not remain excited but re-emits energy aselectromagnetic radiation. This is as a result of an electron 'falling' from a higher

energy level to one of lower energy. This electron transition results in the release

of a photon from the atom of an amount of energy (E = h) equal to the

difference in energy of the electronic energy levels involved in the transition.

In a sample of gaseous hydrogen where there are many trillions of atoms all of

the possible electron transitions from higher to lower energy levels will take

place many times. A prism can now be used to separate the emitted

electromagnetic radiation into its component frequencies (wavelengths orenergies). These are then represented as spectral lines along an increasing

frequency scale to form an atomic emission spectrum.

http://en.wikipedia.org/wiki/Electromagnetic_radiationhttp://en.wikipedia.org/wiki/Electromagnetic_radiationhttp://en.wikipedia.org/wiki/Frequencyhttp://en.wikipedia.org/wiki/Light_emissionhttp://en.wikipedia.org/wiki/Heathttp://en.wikipedia.org/wiki/Electronshttp://en.wikipedia.org/wiki/Electronshttp://en.wikipedia.org/wiki/Chemical_elementhttp://en.wikipedia.org/wiki/Chemical_elementhttp://en.wikipedia.org/wiki/Electromagnetic_radiationhttp://en.wikipedia.org/wiki/Electromagnetic_radiationhttp://en.wikipedia.org/wiki/Frequencyhttp://en.wikipedia.org/wiki/Light_emissionhttp://en.wikipedia.org/wiki/Heathttp://en.wikipedia.org/wiki/Electronshttp://en.wikipedia.org/wiki/Chemical_element


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Lyman series occurs when electrons drop from higher energy levels to the

ground state (n = 1), in this series, the most amount of energy is released and

thus the smallest wavelength and highest frequency. In the Balmer series, all

electrons drop to the n= 2 level, here energies released are not as high as in the

Lyman series. Remember energy is directly proportional to frequency butindirectly proportional

Note that within each series, the spectral lines get closer together with increasing

frequency. This suggests that the electronic energy levels get closer the more

distant they become from the nucleus of the atom.

No two elements have the same atomic emission spectrum; the atomic emission

spectrum of an element is like a fingerprint.


Lyman series occurs when electrons drop

from higher energy levels to the ground

state (n = 1), in this series, the mostamount of energy is released and thus the

smallest wavelength and highest

frequency. In the Balmer series, all

electrons drop to the n= 2 level, here

energies released are not as high as in the

Lyman series. Remember energy is

directly proportional to frequency butindirectly proportional to wavelength.


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Atomic orbitals

Quantum number Definition

First principal quantum number (n) This corresponds to the shell number

e.g. the 1st shell has n=1, 2nd shell has

n=2

1. Orbital volume of space in which there is a 95% chance of finding an

electron

2. Subshell a group of orbitals with the same energy i.e. they are

degenerate

e.g. 3p subshell which has 3 orbitals of the same energy

3. Shell a group of orbitals and/or subshells with the same principal quantumnumber. n =1 shell is called the K shell, n=2, the L shell, n=3 the M shell

etc.

Principal quantum number Types of orbitals/subshells present

in the shell

n=1 1s orbital

n=2 2s orbital and 2p subshell (which

contain THREE 2p orbitals)n=3 3s orbital, 3p subshell (THREE p

orbitals) and 3d subshell (FIVE d

orbitals)

The relative energies of s, p and d orbitals up to principal quantum

number 4

Note: The 4s orbital is LOWER than the 3d orbital. Therefore electrons will

enter the 4s orbital first before the 3d



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Shapes of atomic orbitals (s orbital and p orbital respectively)

Order of filling electrons in orbitals & sub-shells (electronic

configuration)

1s 2s 2p 3s 3p 4s 3d

Each orbital can hold only TWO electrons. Electrons

entering sub-shells

containing 2 or more orbitals enter and occupy the orbitals SINGLY

before pairing. An orbital of lower energy must be filled FIRST beforeelectrons enter an orbital of a higher energy level.

Element 1s 2s 2px 2py 2px 3s 3px 3py 3pz 4s 3d

1H 1

2He 2

3Li 2 1

4Be 2 2

5B 2 2 1

6C 2 2 1 1

7N 2 2 1 1 1

8O 2 2 2 1 1

9F 2 2 2 2 1

10Ne 2 2 2 2 211Na 2 2 2 2 2 1

19K 2 2 2 2 2 2 2 2 2 1

20Ca 2 2 2 2 2 2 2 2 2 2


Remember for each p sub-

shell, there are 3 p orbitals in

x, y and z axis calledpx, py and pz. They are

perpendicular to each other.

They are of the same energy

level and are called

degenerate.


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Element 1s 2s 2px 2py 2pz 3s 3px 3py 3pz 4s 3d

21Sc 2 2 2 2 2 2 2 2 2 2 1

22Ti 2 2 2 2 2 2 2 2 2 2 2

23V 2 2 2 2 2 2 2 2 2 2 324Cr* 2 2 2 2 2 2 2 2 2 1 5

25Mn 2 2 2 2 2 2 2 2 2 2 5

26Fe 2 2 2 2 2 2 2 2 2 2 5

27Co 2 2 2 2 2 2 2 2 2 2 7

28Ni 2 2 2 2 2 2 2 2 2 2 8

29Cu* 2 2 2 2 2 2 2 2 2 1 10

30Zn 2 2 2 2 2 2 2 2 2 2 10

* Indicates that the electronic configuration is not what is expected.

For Cr what would have been expected would be 3d

4 4s2, however, half

filled and totally filled shells/orbitals are very stable and thus more

preferred than any other configuration. Therefore the electrons half fill the

3d subshell with 5 electrons and half fill the 4s orbital with 1 electron.

For Cu, what would be expected was 3d

9 4s2, again the combination of atotally filled subshell and a half filled orbital is more stable than just a filled

orbital and a partly filled subshell. Therefore the electrons adopt the more

stable configuration.

Evidence of discrete energy levels using emission spectra

Based on information given above, it is shown that energy occupy different

orbitals or even-subshells and in essence occupy discrete energy levels.When elements undergo emission spectroscopy and produce an emission

spectrum, a series of lines are shown like the emission spectrum of hydrogen

shown below.



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But how are these lines explained and how do they show evidence of

discrete energy levels?

When the electronsin the elementabsorb energy, they move to higher energylevels and are no longer in the ground state (lowest energy state), they are

now in an excited state (state of higher energy). As the electrons release the

energy absorbed and leave the excited state to return to the ground state, the

excess energy is emitted, the wavelength of which refers to the discrete lines of

the emission spectrum. If discrete energy levels were NOT present, no lines

could EVERbe formed in an emission spectrum.

Emission spectrum of hydrogen

When a gaseous hydrogen atom in its ground state is excited by an input ofenergy, its electron is 'promoted' from the lowest energy level to one of

higher

energy (similar from moving from a lower rung in a ladder to a higher

rung). The atom does not remain excited but re-emits the excess energy as

electromagnetic radiation. This is as a result of an electron 'falling' from a higher

energy level to one of lower energy. This electron transition results in the release

of a photon from the atom of an amount of energy (E = hv) equal to the

difference in energy of the electronic energy levels involved in the transition.

Note E = energy, h = Plancks constant and v = frequency of wavelength ofradiation emitted

In a sample of gaseous hydrogen where there are many trillions of atoms all of

the possible electron transitions from higher to lower energy levels will take

place many times. A prism can now be used to separate the emitted

electromagnetic radiation into its component frequencies (wavelengths or

energies). These are then represented as spectral lines along an increasing

frequency scale to form an atomic emission spectrum.

http://en.wikipedia.org/wiki/Electronshttp://en.wikipedia.org/wiki/Electronshttp://en.wikipedia.org/wiki/Chemical_elementhttp://en.wikipedia.org/wiki/Chemical_elementhttp://en.wikipedia.org/wiki/Electronshttp://en.wikipedia.org/wiki/Chemical_element


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In each spectra a group of lines are see together which is classified as a series.

There are 3 series which are of significance.

The Lyman series occurs when electrons drop from higher energy levels to the

ground state (n = 1), in this series, the most amount of energy is released andthus the smallest wavelength and highest frequency. This is why Lyman series

corresponds to the ultra-violet region (high energy)

The Balmer series occurs when electrons drop from higher energy levels to the

n = 2 level, here energies released are not as high as in the Lyman series. This

corresponds to the visible region of electromagnetic (EM) spectrum.

The Paschen series occurs when electrons drop from higher energy levels to then = 3 level. This corresponds to the infra-red region of electromagnetic (EM)

spectrum.

No two elements have the same atomic emission spectrum; the atomic

emission spectrum of an element is like a fingerprint.

Ionisation energy

It can be quoted more accurately as either 1st, 2nd, 3rd, 4th etc ionisation

energy. For our purposes, we will deal with the 1st ionisation energy.

The 1st ionisation energy is the energy required to remove a mole of

electrons from a mole of gaseous atoms to form a mole of gaseous

univalent ions. A (g) A+ (g) + e-

Trend of 1st ionisation energies

Ionisation energies

generally increase going across a period

Remember two factors must be considered: (1) proton number increases

sequentially going across a period i.e. greater nuclear attraction for the

outermost electron(s) and (2) number of electrons are also increasing.

Although the addition of electrons into the shell causes repulsion and thus

would increase the atomic radius, the predominant factor is the increasedeffective nuclear charge (which is the residual attraction of the nucleus

and the outermost electron(s) after shielding of the inner electrons) i.e. more

energy would be needed to remove the outermost electron(s). Thus

ionisation energy increases from left to right of a period



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Ionisation energies decreases going down a group

Although nuclear charge increases, the dominant factor is the increasing number

of shells between the nucleus and the outermost electron(s). This results in

increased shielding of the nuclear charge, therefore less attraction of the nucleusand the outermost electron(s) i.e. less energy needed to remove an outermost

electron. Thus ionisation energy decreases down a group.

Atypical behaviour seen in period 3 for Mg & Al AND P & S (period 3)

In period 3, Mg has E.C. of [Ne] 3s2, while Al has E.C. of [Ne] 3s2 3px1, in Al

the outermost electron (3p) is at a higher energy level than the outermost

electron in Mg (3s), therefore less energy is needed to remove it. Or using a

different explanation, the valence electron in Al experiences more shielding i.e.less nuclear attraction than one of the valence electrons in Mg i.e. less energy

needed to remove it from Al than for Mg.

In period 3 for P and S, the explanation needed is somewhat different. For P,

the E.C. is [Ne] 3s2 3p3, while for S the E.C. is [Ne] 3s2 3p4. In the 3p subshell of

P, the half-filled subshell represents a very stable configuration since it

represents a system of minimum repulsion as each electron occupies one orbital

singly. A lot of energy would be needed to disrupt this configuration. While in

S, 3p subshell experiences electron-electron repulsion in one of its orbital whichraises the energy of the system, therefore it is LESS stable and LESS energy

would be needed to remove one of the valence electrons.

Below is a diagram for the 1st ionisation energy of period 3 elements. Note

the circles show the areas of atypical behaviour



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Evidence of sub-shells using ionization data

The graph below shows the successive ionization energies for an atom of

sodium:

The electronic structure for sodium is 1s2 2s2 2p6 3s1.

The energy required to

remove the first electron is relatively low. This corresponds to the loss ofone 3s electron. To remove the second electron needs a much greater energy

because this electron is closer to the nucleus in a 2p orbital. There is a steady

increase in energy required as electrons are removed from 2p and then 2s

orbitals.

The removal of the tenth and eleventh electrons requires much greater

amounts of energy, because these electrons are closer to the nucleus in the 1s

orbital.

Large jumps in energy shown by the circles, indicate moving from one

principal quantum number to another. The smaller, more gradual increases

indicate going moving within subshells as the energies of the electrons will

slowly decrease resulting in more and more energy needed to remove them.



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How to derive group number of an element from successive ionization

energies

A large jump (usually an increase of 3 or more times the amount) between

two successive ionisation energies is typical of suddenly breaking in to an

inner level. You can use this to work out which group of the Periodic Table

an element is in from its successive ionisation energies.

Example 1 Magnesium (1s22s22p63s2) is shown with the following

successive ionisation energies:

Here the big jump occurs after the second ionisation energy. It means that

there are 2 electrons which are relatively easy to remove

(the 3s2 electrons),

while the third one is much more difficult (because it comes from an inner

level - closer to the nucleus and with less screening). Mg is therefore ingroup II

Example 2 Silicon (1s22s22p63s23px13py

1) is shown with the following

successive ionisation energies:

Here the big jump comes after the fourth electron has been removed. The

first 4 electrons are coming from the 3rd shell orbitals; the fifth from. Silicon

is therefore in group IV

To try on your own

Decide which group an atom is in if it has the following successive

ionisation energies:

END OF ATOMIC STRUCTURE



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Worksheet

1.Write the electronic configurations of the following atoms or ions

a) 20Ca..

b) 7N3-..

c) 26Fe2+..

d)29Cu..

2.



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3.



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