Atomic Structure

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
Daltons Atomic Theory

• The Beginning of modern atomic theory is credited
  to John Dalton.

5
Daltons Atomic Theory

• Daltons 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
Daltons 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
Daltons Atomic Theory

• We know now that Daltons 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, Daltons 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 1800s 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 1800s
  early 1900s.
• 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 -1.75881 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 atoms
  structure.

21
Rutherford and the Nuclear Atom

• If Thompsons 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

• Rutherfords 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 atoms mass is in the nucleus.
• The nuclear diameter is 1/10,000 to 1/100,000
  times less than atoms 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 elements box.

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
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

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.

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 2.331411
• Mass of 12C atom
• From this mass ratio, the isotopic mass of the
  28Si can be found
• 2.331411 x 12 amu 27.97693 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
  10.811 amu. The masses of the two naturally
  occurring isotopes are 510B and 511B, are 10.013
  and 11.009 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 0.199 and abundance of 10B
  19.9.
• The multiplier for 11B is (1-x) thus the fraction
  of 11B is 1-0.199 0.801 and the abundance of
  11B is 80.1.

45
Home Work

• 1. Calculations Chemistry 9th Edition, Chapter
  4, Exercises 28 38.
• 2. What are the main points in Daltons 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.