1. 1
Atomic Structure

2. Introduction
All substances are made up of matter and the fundamental unit of matter is the atom.
The atom constitutes the smallest particle of an element which can take part in chemical reactions and may or may not exist independently.

3. Introduction
Most of what is known about atomic structure is based on basically 2 types of research:-
Electrical nature of matter
Interaction of matter with light energy

4. Dalton’s Atomic Theory
The Beginning of modern atomic theory is credited to John Dalton.

5. Dalton’s Atomic Theory
Dalton’s model was formulated on a number “laws” about how matter behaves in a chemical reaction.
These laws were based on experimental evidence
Observations on many chemical substances & their reactions
These laws are the foundation on which the modern atomic theory is based

6. Dalton’s Atomic Theory
Dalton formulated his theory:
All matter is made up of atoms (small, indivisible, indestructible, fundamental particles)
Atoms can neither be created or destroyed (they persist unchanged for all eternity)
Atoms of a particular element are all alike (in size, mass & properties)
Atoms of different elements are different from one another (different sizes, masses & properties)
A chemical reaction involves either the union or the separation of individual atoms

7. Dalton’s Atomic Theory
We know now that Dalton’s theory is not entirely true, for example:
Atoms are not the most fundamental particles – they are composed of smaller particles
Atoms can be created or destroyed but a nuclear process is needed to do so
Nonetheless, Dalton’s model was superb for his time and it laid the foundation for further developments in atomic theory.

8. Fundamental Particles
Three fundamental particles make up atoms.
The following table lists these particles together with their masses and their charges.

9. The Discovery of Electrons

10. The Discovery of Electrons
Earliest evidence for atomic structure was supplied in the early 1800’s by the English chemist, Humphrey Davy.
Davy passed electricity through compounds and noted:
that the compounds decomposed into elements.
concluded that compounds are held together by electrical forces.

11. The Discovery of Electrons
Most convincing evidence came from Cathode Ray Tubes experiments performed in the late 1800’s & early 1900’s.
Consist of two electrodes sealed in a glass tube containing a gas at very low pressure.
When a voltage is applied to the cathodes a glow discharge is emitted.

12. The Discovery of Electrons
These “rays” are emitted from cathode (- ve end) and travel to anode (+ve end).
Cathode Rays must be negatively charged!

13. The Discovery of Electrons
J.J. Thomson modified the cathode ray tube experiments in 1897 by adding two adjustable voltage electrodes.
Studied the amount that the cathode ray beam was deflected by additional electric field.

14. The Discovery of Electrons
Modifications to the basic cathode ray tube experiment show the nature of cathode rays
(a) A cathode ray discharge tube, showing the production of a beam of electrons (cathode rays). The beam is detected by observing the glow on a flourescent screen.
(b) A small object placed in front of the beam, casts a shadow indicating that cathode rays travel in straight lines.

15. (c) Cathode rays have a negative electrical charge, as demonstrated by their deflection in an electrical field.
(d) Interaction with a magnetic field also consistent with negative charge.
(e) Cathode rays have mass, as shown by their ability to turn a small paddle wheel in their path.

16. The Discovery of Electrons
Thomson used his modification to measure the charge to mass ratio of electrons.
Charge to mass ratio
e/m = x 108 coulomb/g of e-
Thomson named the cathode rays electrons.
Thomson is considered to be the “discoverer of electrons”.

17. Canal Rays and Protons
In 1886 Eugene Goldstein noted that cathode ray tube also generated streams of positively charged particles that moved toward the cathode.

18. Canal Rays and Protons
Particles move in opposite direction of cathode rays.
Called “Canal Rays” because they passed through holes (channels or canals) drilled through the negative electrode.
- Canal rays must be positive.
Goldstein postulated the existence of a positive fundamental particle called the “proton”.

19. Model for Atomic Structure
By early 1900s it was clear that atoms contained regions of +ve and -ve charge.
But how these charges were distributed was still unclear.
1st model for the structure of the atom was proposed by Thompson based on the following:
Atoms contain small –ve charged particles (electrons)
Atoms of an element behave as if they had no electrical charge
So there must be something in the atom to neutralize the –ve electrons (protons not yet discovered)

20. Rutherford and the Nuclear Atom
Further insight into atomic structure was provided by Ernest Rutherford.
He has established that - particles were +ve charged particles
They are emitted by some radioactive atoms (when they disintegrate spontaneously)
Bombarded thin Au foils with - particles from a radioactive source
Gave us the basic picture of the atom’s structure.

21. Rutherford and the Nuclear Atom
If Thompson’s model was correct then any - particles passing through the foil would be deflected by small angles.
Unexpectedly most of the - particles passed through the foil with little or no deflections (shown in black).
Many were deflected through moderate angles (shown in red).
These deflections were surprising, but the 0.001% of the total that were reflected at acute angles (shown in blue) were totally unexpected!

22. Rutherford and the Nuclear Atom

23. Rutherford and the Nuclear Atom
Rutherford’s major conclusions from the -particle scattering experiment:
The atom is mostly empty space.
It contains a very small, dense center called the nucleus.
Nearly all of the atom’s mass is in the nucleus.
The nuclear diameter is 1/10,000 to 1/100,000 times less than atom’s radius.

24. Neutrons
James Chadwick in 1932 analyzed the results of -particle scattering on thin Be films.
Chadwick recognized existence of massive neutral particles which he called neutrons.
Chadwick discovered the neutron.

25. Atomic Number The atomic number = # of protons in the nucleus.
Sometimes given the symbol Z.
On the periodic table Z is the uppermost number in each element’s box.
Mass
number
Charge of particle
Symbol of
the atom
(= p + n)
(= # of p )
Atomic
number

26. Atomic Number
In 1913 H.G.J. Moseley realized that the atomic number determines the element.
The elements differ from each other by the number of protons in the nucleus.
So… it is the number of protons that determine the identity of an element
The number of electrons in a neutral atom is also equal to the atomic number.

27. Nucleon Number and Isotopes
Mass
number
Charge of particle
Symbol of
the atom
(= p + n)
(= # of p)
Atomic
number
Nucleon number (formerly Mass number) is given the symbol A.
A = # of protons + # of neutrons.
If Z = proton number and N = neutron number
Then A = Z + N

28. So, the Standard Notation used to show mass and proton numbers is:
Charge of particle
Symbol of
the atom
(= p + n)
(= # of p )
Atomic
number
Can be shortened to this symbolism.

29. Mass Number and Isotopes
Isotopes are atoms of the same element but with different neutron numbers.
Isotopes have different masses and A values but are the same element.
Protium
(Hydrogen - 1)
Deuterium
(Hydrogen - 2)
Tritium
(Hydrogen - 3)
Isotopes of Hydrogen

30. Mass Number and Isotopes
The stable oxygen isotopes provide another example.
1. 16O is the most abundant stable O isotope. How many protons and neutrons are in 16O?
2. 17O is the least abundant stable O isotope. How many protons and neutrons are in 17O?
3. 18O is the second most abundant stable O isotope. How many protons and neutrons in 18O?

31. Mass Spectrometry & Isotopic Abundances
Identifies chemical composition of a compound or sample on the basis of the mass-to-charge ratio of charged particles.
A gas sample at low pressure is bombarded with high-energy electrons.
This causes electrons to be ejected from some of the gas molecules  creating +ve ions.
Positive ions then focused into a very narrow beam and accelerated by an electric field.
Then passes through a magnetic field which deflects the ions from their straight path.

32. Mass Spectrometry
There are four factors which determine the extent of deflection:
Accelerating voltage
Higher voltages  beams move more rapidly and deflected less than slower moving beams produced by lower voltages.
Magnetic field strength
Stronger fields give more deflection
Masses of particles
Heavier particles deflected less than lighter ones
Charge on particles
Particles with higher charges interact more strongly with magnetic fields and are thus deflected more than particles of equal mass with small charge.

33. Mass Spectrometry Mass Spectrometry A modern mass spectrometer
Fig. 5-10a, p. 176

34. Mass Spectrometry & Isotopic Abundances
Mass spectrum of Ne+ ions shown below.
How do scientists determine the masses and abundances of the isotopes of an element?
Neon consists of 3 isotopes, of which Neon-20 is the most abundant (90.48%).
The number by each peak corresponds to the fraction of all the Ne+ ions represented by the isotope with that mass.

35. Mass Spectrometry & Isotopic Abundances
The mass of an atom is measured relative to the C-12 atom
Its’ mass is defined as exactly 12 atomic mass units (amu)
Therefore the amu is 1/12 the mass of a C-12 atom
Example: What is the mass in amu of a 28Si atom?
The spectrometer will measure the ratio of the mass of an 28Si atom to 12C:
Mass of 28Si atom =
Mass of 12C atom
From this mass ratio, the isotopic mass of the 28Si can be found:
x 12 amu = amu
The mass of the isotope relative to the mass of the C-12 isotope

36. Table 5-3, p. 178

37. Isotopes Small differences in physical properties.
Similar chemical properties because isotopes have same number of p and e.
Some isotopes are radioactive.
nuclear behavior of isotopes is unique
Radioactive isotopes are biologically useful
Example: radioactive I-131 to study thyroid gland

38. Atomic Weight
The relative atomic weight (also called relative atomic mass) of an element is the weighted average of the masses of its stable isotopes.

39. Atomic Weight Atoms are amazingly small.
Their masses are compared with the mass of an atom of the carbon-12 isotope, as the standard.
One atom of the C-12 iostope weight exactly 12 units (Atomic mass units, amu).
E.g. an atom of the most common isotope of Mg weighs twice as much as one atom of C-12, its relative isotopic mass is 24.

40. Atomic Weight “weighted average” e.g. Chlorine Cl- 35 75% Cl-37  25%
If you had 100 atoms, 75 would be Cl-35 and 25 would be Cl-37.
The weighted average is closer to 35 than 37 because there are more Cl-35 than Cl-37 atoms.

41. Atomic Weight Example: Naturally occurring Cu consists of 2 isotopes.
It is 69.1% 63Cu with a mass of 62.9 amu,
and 30.9% 65Cu, which has a mass of 64.9 amu.
Calculate the atomic weight of Cu to one decimal place.

42. Atomic Weight
Example: The relative atomic mass of boron is amu. The masses of the two naturally occurring isotopes are 510B and 511B, are and amu, respectively. Calculate the fraction and percentage of each isotope.
You do it!
This problem requires a little algebra.
A hint for this problem is x + (1-x) = 1

43. Atomic Weight

44. Atomic Weight
Note that because x is the multiplier for the 10B isotope, our solution gives us the fraction of natural B that is 10B.
Fraction of 10B = and % abundance of 10B = 19.9%.
The multiplier for 11B is (1-x) thus the fraction of 11B is = and the % abundance of 11B is 80.1%.

45. Home Work
1. Calculations: Chemistry 9th Edition, Chapter 4, Exercises 28 – What are the main points in Dalton’s atomic theory? 3. Briefly outline how the mass spectrometer works to help determine the isotopic abundance and isotopic mass. Include a diagram in your answer.