1791383751

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Atom 1

ATOMIC STRUCTURE

An atom is the smallest part of an element which can ever exist.

Atoms consists of a small central region called nucleus in which all the mass of the atom (protons and neutrons) are concentrated.

This nucleus is surrounded by a much larger volume in which the electrons move.

The nucleus of an atom is about one ten-thousandth of the size of the whole atom. If we magnified an atom to the size of a football stadium, the

nucleus would be represented by a pea placed at the centre of the pitch.

Particle Symbol Charge / C Relative charge Mass / g Relative mass

Proton p+ +1.602 x 10

-19 +1 1.673 x 10

-24 1.007 3

Neutron n 0 0 1.675 x 10-24 1.008 7

Electron e- -1.602 x 10-19 -1 9.109 x 10-28 5.4 x 10-4

(negligible)

Any nucleus is characterized by two numbers :

Z A X where Z is the atomic number = no. of protons

A is the mass number = no. of protons + neutrons

Atomic number is more significant because :

1 all elements have their own unique atomic numbers

2 all atoms of the same element must have the same number of protons but their mass may be different from one another as a

result of the existence of isotopes (atoms of the same element with different number of neutrons)

Therefore, elements in the Periodic Table are arranged in the order of atomic numbers.

Relative atomic mass is the ratio of the weighted mean isotopic mass of the atom to121 of the mass of a 12C atom.

At first, the element hydrogen was chosen as the standard against which the masses of other atoms were compared. It was because chemists

realized that it had the smallest atoms which could conveniently be assigned a relative atomic mass of one. Later, when relative atomic masses

could be obtained more accurately, carbon-12 was chosen as the new standard because :

1 carbon is a very common element;

2 being a solid, it is easier to store and transport than hydrogen, which is a gas.

Class Work

Calculate the relative atomic masses for argon and potassium using the following data :

Ar : 36 (0.34) ; 38 (0.06) ; 40 (99.6)

K : 39 (93.12) ; 40 (0.12) ; 41 (6.76)

MASS SPECTROMETER

Relative masses of atoms and of molecules can be accurately

determined by a mass spectrometer.

An analogy to the mass spectrometer would be to roll a bowling ball and

a basketball at the same speed at a target while a stiff crosswind isblowing. The basketball is lighter, therefore, its path is more readily

changed by the crosswind.

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Atom 2

component function

A to vaporize the sample (which may be an element or a compound)

B

the sample is bombarded by fast moving electrons to form positive ions

X(g) + e

-

→

X

+

(g) + 2 e

-

fast slow

C to accelerate the beam of positive ions

D

to deflect the ion beam so that ions of a particular mass/charge (m/e) ratio

are focused into the ion-detector

(the lighter the positive ions, the greater the deflection)

E to detect the signal and pass it on to a recorder

By varying the strength of accelerating electric field ordeflecting magnetic field, ions of any m/e ratio can be

brought to the ion detector. A mass spectrum showing

the m/e ratio of the ions and the corresponding intensity

(i.e. the relative amount of that ion) can then be traced

out by a recorder. In most mass spectra, the values of

the m/e ratio can be converted to the relative masses of

the particles if the charges on the ions are taken to be

one.

Each peak in the mass spectrum represents the relative

abundance of a particular type of particle with a certain

isotopic mass.

(No instrumental details and mathematical treatments are

required)

Class Work

The following mass spectrometer trace was obtained for a naturally occurring sample of an element X .

80

69.95

60Relative

abundance 40

30.05

20

0 203 205Relative isotopic mass

(i) Give an interpretation of this trace.

(ii) Calculate the relative atomic mass of the naturally occurring sample to 4 sig. fig.

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Atom 3 Class Work

The mass spectrum of chlorine molecule contained three peaks at m/e 70, 72 and 74 of relative intensity 9 : 6 : 1 respectively.

Deduce the relative atomic mass of chlorine.

The material to be analyzed may be an element or a compound :

1 when sample is an element, the peaks in the mass spectrum can give information about various isotopes of the element

2 when sample is a compound, the peak with the highest m/e ratio will most likely correspond to the ‘molecular ion’, i.e. the

molecule which has lost only a single electron

Class Work

Sketch the mass spectrum for (i) methane, CH4 ; and (ii) hydrogen chloride, HCl.(Assuming 1H, 12C and 14N have 100% abundance among their isotopes)

ELECTRONIC STRUCTURE OF ATOMS

Our knowledge of the way in which electrons are

distributed in an atom comes from two sources of

evidence :

1 study of emission spectrum of atomic hydrogen

2 study of ionization energies of the elements

Emission Hydrogen Spectrum

When an electrical discharge is passed through

hydrogen at low pressure, a coloured glow is observed.

Upon dispersion by a prism, the radiation can be

spread out into a few sharp, coloured lines, which are

commonly known as the emission hydrogen

spectrum.

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Atom 4 Characteristics of Emission Hydrogen Spectrum

1 It consists of several series of discrete lines in different parts of the electromagnetic spectrum; e.g. Lyman series in the

ultraviolet (UV) region, Balmer series in the visible region and Paschen series in the infra-red (IR) region

2 It consists of coloured lines against a dark background

3 The lines in each series converge (get closer) towards higher frequencies (or shorter wavelengths), until they eventually

merge into a continuum

4 Every element has its own unique pattern of lines in its emission spectrum

4 5 6 7 8frequency, ν / 1014 Hz

Balmer worked out a relationship to describe the pattern of lines in the visible spectrum :

1

n

2

ν α

The lines in all series are found to fit in the general formula :

where λ is the wavelength of the particular line

R is Rydberg’s constant ( = 1.097 x 107 m

-1 )

n, m are integers which, in turn, have the following set of values :

Lyman series n = 1 m = 2, 3, 4, 5 . . . ∞ in the UV region

Balmer series n = 2 m = 3, 4, 5, 6 . . . ∞ in the visible region

Paschen series n = 3 m = 4, 5, 6, 7 . . . ∞ in the IR region

Class Work

The wavelength of the first line in the Lyman series is found to be 1.215 x 10-7

m.

(a) Calculate the wavelength of the third line in the Balmer series.

(b) Sketch the first five lines in the Balmer series in the above emission spectrum.

1st line 2nd line 3rd line 4th line 5th line

m

1

λ

λ / 10-7 m

ν / 1014 Hz

= R (1

n -

1

m )

2 2

1

λ

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Atom 5 Interpretation of Emission Hydrogen Spectrum

To account for the hydrogen spectrum, it is assumed that electrons in an atom is quantized (i.e. electrons can only move in

orbits with a certain fixed amount of energy). Under normal conditions, the electrons in an atom or ion fill the lowest energy

levels first (the ground state). Upon absorption of sufficient energy (e.g. via electrical discharge), it is possible to promote (excite)

an electron from a lower energy level to a higher one (the excited state).

The electron is now unstable in the higher energy level, so it

will emit the excess energy as radiation and drop back into thelower energy level. The energy difference between the higher

and lower energy levels can have only certain fixed values

because the energy levels themselves are fixed. This means

that transition of an electron from a higher to a lower level

emits a photon (discrete amount of energy in the form of

radiation) with energy equal to the difference in energy

between the two energy states, hence produces a discrete

line in the emission spectrum.

The spectrum lines in each series converge towards higher

frequencies as a result of the fact that the energy difference

between orbits decrease with increase in principal quantum

number (i.e. higher energy levels).

The relationship between the energy ( E ) of radiation and its

frequency ( ν) is

E = h ν νν ν

where h is Planck’s constant ( = 6.63 x 10-34

J s)

The atoms of each element have a unique arrangement of electrons with definite energy levels. When the atom is excited,

electrons are brought to higher energy levels. As the electrons return to lower energy levels, the atom emits radiation of definite

wavelengths. This results in a unique emission spectrum which can provide useful information about the atom and serve as an

identification for the element concerned.

Ionization Energies

If sufficient energy is applied to an atom to excite an electron from its ground state to just beyond the highest possible one, then the

electron can escape from the atom. The atom is said to be ionized.

In an atom, the highest possible energy level corresponds to the frequency at which the lines in the spectrum come together. So, by

determining the frequency at which the converging spectral lines come together, we can find the ionization energy of an element.

This particular frequency is called the convergence limit.

Class Work The wavenumber of the first line in the Balmer series is found to be 1.525 x 10 6 m-1. Calculate the ionization energy per mole of

hydrogen atoms.

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Atom 6 Successive Ionization Energy - Evidence of Shells

If an atom containing several electrons is provided with sufficient energy it will lose one electron. Additional supplies of energy

will result in the removal of a second electron, then a third, a fourth, and so on, i.e. a succession of ionizations is possible.

Class Work

The successive ionization energies of beryllium are 900, 1760, 14900 and 21060 kJ mol-1 respectively.

(i) Write equations to represent the succession of ionization energiesof beryllium.

(ii) Sketch the trend for the succession ofionization energies of beryllium.

I.E.

1 2 3 4

Each successive ionization requires a progressively larger amount of energy than the preceding one because :

1 every time an electron is being removed, the resulting particle carries one more positive charge and exerts a stronger

electrostatic attraction on that electron

2 electrons are being removed from a progressively lower energy levels, which are much closer to the nucleus and thus being

held more tightly

A plot of succession of ionization of an element gives information about its electronic arrangement (e.g. 2, 2 in the case of Be). It

also proves that two of the four electrons spend most of the time closer to the nucleus than the other two. They are said to occupy

different electron shells.

Class Work Sketch the successive ionization energies of sodium and carbon.

Evidence of Sub-shells

1st I.E.

2 4 6 8 10 12 14 16 18 20atomic number

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Atom 7 Trends in the plot of 1st I.E. of the first twenty elements :

1 Down the group, I.E. decrease

one more shell

distance between nucleus & outermost electron increase

more readily to remove(ionize) outermost electron

2 Across the period, I.E. increase

size decrease

outermost electron being held by nucleus more tightly

more difficult to remove outermost electron

3 Within each period, I.E. is minimum at Group I and maximum at Group 0

atoms tend to have their outermost shell full-filled in order to attain an extra stability

4 Within each period, broken trend is found in a pattern of 2, 3, 3

suggest the presence of sub-shells based on similar reason as in (3)

Electronic Configurations

Atomic Orbitals

Louis de Broglie postulated that matter could have both particle and wave properties. This postulate was later confirmed by the experiment thatdemonstrated the diffracting property of electrons, proving that electrons did indeed possess wave properties.

Electrons, therefore, are not localized in any fixed orbits. We cannot know exactly where an electron is at any particular moment,

but we can only describe the probability of finding the electron in a certain position at any time.

The region in space where an electron is likely to be found is called an orbital, and is best illustrated by a

charge cloud diagram where the dot density represents the probability of finding the electron.

Suppose a photograph of the electron of hydrogen could be taken at any instant, and a second photograph taken at

an instant later when that electron would occupy a different position. Millions of such photographs superimposed

would generate an electron density diagram, the best representation of an atomic orbital.

Atomic orbitals have no definite boundary owing to the probability description. Instead, their shapes are more significant

properties in describing atomic orbitals :

s-orbital -- spherically symmetric (i.e. variation is exactly the same in any direction from the nucleus)

-- electron cloud density is densest near the centre, but falls away as the distance increases from the nucleus

-- electron density is zero at the nucleus and at infinite distance from the nucleus

p-orbital -- dumbbell-shaped

-- each p-orbital (namely px, py and pz) has its axis mutually perpendicular to each other

d -orbital -- 4 out of 5 orbitals have 4 lobes extending out perpendicular to each other

-- the last one has 2 lobes extending along z-axis with a doughnut-shaped ring around the centre of x-y plane

The average distance of s electrons from the nucleus is

less than that of p electron, so that s electrons are more

firmly held by the nucleus.

p x p y p z

d xy d xz d yz

d x 2 – y 2 d z 2

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Atom 8 Building Up Electronic Configurations

Rules in building up electronic configurations :

1 Aufbau’s building up principle – electrons will enter

the orbitals in order of increasing energy

2 Hund’s rule – orbitals of the same energy must be

occupied singly before pairing occurs

3 Pauli’s exclusion principle – no orbitals can

accommodate more than TWO electrons

Notation in writing electronic configurations :

no. of electrons

main shell 1 s 2 in the orbital

type of sub-shells / orbitals

Electronic Configurations from H to Kr

atomic

number element symbol notations using 1s, 2s, 2p, … etc. ‘electron-in-boxes’ diagram

1

2

hydrogen

helium

H

He

1s1

1s2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

19

20

n = 1

4d

3d

4s

4 p

2 p

3s

2s

1s

3 p

n = 3

n = 2

n = 4

quantum shells sub-

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Atom 9

atomic

number element symbol notations using 1s, 2s, 2p, … etc. ‘electron-in-boxes’ diagram

21

22

23

24

25

26

27

28

29

30

31

32

33

34

35

36

Some important points to note :

1 There is a division of the Periodic Table into four different areas, namely

the s-, p-, d - and f - blocks, depending on the type of the outermost electron

2 4s sub-shell always has a top priority than 3d sub-shell in filling up or

removing electrons

3 Chromium and copper exhibit abnormal configurations in order to attain an

extra stability from the half-filled and full-filled 3d sub-shell respectively

1s

2s 2 p

3s 3 p 3d

4s 4 p 4d 4 f

5s 5 p 5d 5 f

6s 6 p 6d 6 f

Summary of Electronic Structure of Atoms

A study of :

1 emission spectrum of atomic hydrogen proves that electrons can exist in a certain discrete energy levels

2 succession of ionization for a particular element leads to its electronic configuration

3 the trend of first ionization energies introduces the presence of sub-shells within main shells

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