CHAPTER 5

The Structure
of Atoms
1
Fundamental Particles

Three fundamental particles make up atoms:
Particle
Mass (amu)
Charge
Electron (e-)
0.00054858
-1
Proton (p,p+)
1.0073
+1
Neutron(n,n0)
1.0087
0
2
The Discovery of Electrons

Late 1800’s & early 1900’s
Cathode ray tube experiments showed that
very small negatively charged particles are
emitted by the cathode material.

1897 – J. J. Thomson
Modified the cathode ray tube and
measured the charge to mass ratio of
these particles. He called them electrons.
(Nobel prize in physics, 1906)
3
The Discovery of Electrons


1909 – Robert A. Millikan
Determined the charge and the mass of
the electron from the oil drop experiment.
(The second American to win
Nobel prize in physics in 1923)
1910 – Ernest Rutherford
Gave the first basically correct picture
of the atom’s structure.
(Nobel prize in chemistry in 1908)
4
Rutherford’s Atom
 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
5
The Discovery of Protons

1913 – H.G.J. Moseley
Realized that the atomic number
defines the element:
 Each element differs from the
preceding element by having one more
positive charge in its nucleus

Along with a number of observations made
by Rutherford and some other physicists,
this led to the discovery of the proton
 The elements differ from each other by
the number of protons in the nucleus
6
The Discovery of Neutrons

1932 – James Chadwick
recognized existence of massive neutral
particles which he called neutrons
(Nobel prize in physics in 1935)
 The atomic mass of an element is mainly
determined by the total number of
protons and neutrons in the nucleus
 The atomic number of an element is
determined by the total number of
protons in the nucleus
7
Mass Number and Atomic Number


Mass number – A
Atomic number – Z




Z = # protons
A = # protons + # neutrons
# protons = # electrons
The way we typically write this:
37
17
Cl
full nuclide symbol
37
E
A
Z
Cl
short nuclide symbol
8
Isotopes

Atoms of the same element but with
different masses


The same element means that the
number of protons is the same,
then different masses mean that
the number of neutrons differs
H
1
1
protium
(or hydrogen)
H
2
1
deuterium
H
3
1
tritium
9
Isotopes: Example
U
O
235
O
238
16
17
U
O
18
10
Experimental Detection of Isotopes

1919-1920 – Francis Aston
Designed the first mass-spectrometer
(Nobel prize in chemistry in 1922)

Factors which determine a particle’s
path in the mass spectrometer:
 accelerating voltage, V
 magnetic field strength, H
 mass of the particle, m
 charge on the particle, q
11
Mass Spectrometry & Isotopes

Mass spectrum of Ne+ ions

This is how scientists determine the masses
and abundances of the isotopes of an element
12
Mass Spectrometry & Isotopes

Let’s calculate the atomic mass of Ne
using the mass-spectrometry data
13
Atomic Weight Scale


A unit of atomic mass (atomic mass unit) was
defined as exactly 1/12 of the mass of a 12C atom
Two important consequences of such scale choice:



The atomic mass of 12C equals 12 a.m.u.
1 a.m.u. is approximately the mass of one atom
of 1H, the lightest isotope of the element with
the lowest mass.
The atomic weight of an element is the weighted
average of the masses of its isotopes
14
Isotopes and Atomic Weight

Naturally occurring chromium consists
of four isotopes. It is
4.31% 50Cr, mass = 49.946 amu
83.76% 52Cr, mass = 51.941 amu
9.55% 53Cr, mass = 52.941 amu
2.38% 54Cr, mass = 53.939 amu
Calculate the atomic weight of chromium
15
Isotopes and Atomic Weight


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.
A.W.(Cu) = (62.9 amu  0.691) + ( 64.9 amu  0.309) =
= 63.5 amu
16
Electromagnetic Radiation

Any wave is characterized by 2 parameters:


Wavelength () is the distance between two
identical points of adjacent waves, for
example between their crests
It is measured in units of distance (m, cm, Å)
Frequency () is the number of wave crests
passing a given point per unit time (for
example, per second)
It is measured in units of 1/time, usually s-1
1 s-1 = 1 Hz (Hertz)
17
Electromagnetic Radiation

The speed at which the wave propagates:
c=

The speed of electromagnetic waves in
vacuum has a constant value:
c = 3.00108 m/s


This is the speed of light
Given the frequency of the
electromagnetic radiation, we can
calculate its wavelength, and vice versa
18
Electromagnetic Radiation
Max Planck
(Nobel prize in physics in 1918)




Electromagnetic radiation can also be
described in terms of “particles” called
photons
Each photon is a particular amount of
energy carried by the wave
Planck’s equation relates the energy of
the photon to the frequency of radiation:
E=h
(h is a Planck’s constant, 6.626·10-34 J·s)
19
Electromagnetic Radiation

What is the energy of green
light of wavelength 5200 Å?
20