Atomic Structure page 1 of 18

ATOMIC STRUCTURE AND THE PERIODIC TABLE

Theoretical change with respect to Dalton’s atomic theory

1. In 1803, atomic theory was revived by John Dalton

a) 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 split 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

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

kg

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Terminology

Term Definition

Atomic/proton number Number of protons in a nucleus of an atom

Nucleon/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 and 9

4B 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 properties

Relative 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

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 of radioactive decay in which an atomic nucleus emits 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 number 4 less and atomic number 2 less.

For example:

although this is typically written as:

Beta decay is a type of radioactive decay in which a fast moving electrons is emitted. The new atom has no change in mass number but an atomic number

increases by 1.

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Gamma rays or gamma-ray (denoted as γ) are forms of electromagnetic

radiation (EMR) or light emissions of a specific frequency produced from sub-

atomic particle interaction, such as electron-positron annihilation and

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

atom.

Band of stability (n/p ratio)

Uses of radioisotopes

1. radiocarbon dating 2. smoke detectors

3. pacemakers 4. medical uses i.e. tracers or chemotherapy

5. irradiation in pest control

Calculations of relative atomic mass from isotopic data

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

an element)

e.g. what is the relative atomic mass of zirconium (Zr) Zr-90 51.5% ,

Zr-91 11.2%, Zr-92 17.1%, Zr-94 17.4 % and Zr-96 2.8%

Ar Zr = (51.5x 90) + (11.2x91) + (17.1x92) + (17.4x94) + (2.8x96) = 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.

Most elements have isotopes. For stable isotopes, 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 the Nuclear Belt of Stability. The plot indicates that lighter nuclides (isotopes) are most stable when the neutron/proton ratio is 1/1. This is the case with any nucleus that has up to 20 protons.

As the atomic number increases beyond 20, a different trend becomes apparent. In this range, it appears that a stable nucleus is able to accommodate more neutrons. Stable isotopes have a higher neutron to proton ratio, rising to 1.5/1 for elements having atomic

numbers between 20 and 83.

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Terminology

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 quantum

number. n =1 shell sometimes shells can be the K shell where n=1, the L

shell where n=2, the M shell where n=3, 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

Remember for each p sub-

shell, there are 3 p orbitals in

x, y and z axis called

px, py and pz. They are

perpendicular to each other.

They are of the same energy

level and are called

“degenerate”.

Due to Pauli’s Exclusion

Principle (no two electrons can have the same 4 quantum states),

this implies that each orbital can

hold only TWO electrons, one

spinning “up” and the other spinning “down”. The Aufbau

Principle states that electrons enter

and fill orbitals of lower energy levels before going to higher

energy levels. Hund’s Rule states

that electrons entering sub-shells

containing 2 or more orbitals must enter and occupy the orbitals

SINGLY before pairing.

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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 2

11Na 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

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 3

24Cr* 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 3d4 4s

2, 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 4s

2, again the combination of a

totally 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.

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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.

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

discrete energy levels?

When the electrons in the element absorb energy, they move to higher energy

levels 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 EVER be formed in an emission spectrum.

Emission spectrum of hydrogen When a gaseous hydrogen atom in its ground state is excited by an input of

energy, 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 = Planck’s constant and v = frequency of wavelength of

radiation 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.

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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 and

thus 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 the

n = 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 1

st, 2

nd, 3

rd, 4

th 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 1

st 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 increased

“effective 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 nucleus

and 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 3px

1, 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 which

raises 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 2s

2 2p

6 3s

1. The energy required to

remove the first electron is relatively low. This corresponds to the loss of

one 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 (1s22s

22p

63s

2) 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 in

group II

Example 2 Silicon (1s22s

22p

63s

23px

13py

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) 7N

3-……………………………..

c) 26Fe2+

………………………..

d)29Cu…………………………..

2.

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3.

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4.Below are incomplete nuclear equations, fill in any missing

information below i.e. any atomic and/or mass numbers missing as

well as any symbols of atoms that were not included.

5.

………………………………………………………………………

…………………..…………………………………………………

…………………………………………..…………………………

………………………………………………………………….

………………………………………………………………………

…………………….………………………………………………

…………………………………………….………………………

…………………………………………………………………….

………………………………………………………………………

……………………..………………………………………………

……………………………………………..………………………

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6. Using the same atoms from question 1, use the orbital diagrams

to illustrate the electronic configurations.

a) Mg

b) C

c) Ne

d) Li

e) K

f) Sc

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g) Fe

h) Co

i) Cu

j) Mn