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Atomic Structure

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. 2

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 3

Dalton’s Atomic Theory

• Beginning of modern atomic theory credited to John Dalton.

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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 5

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 6

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 7

Fundamental Particles

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

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The Discovery of Electrons

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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.

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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.

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The Discovery of Electrons

• These “rays” are emitted from cathode (- ve end) and travel to anode (+ve end).

– Cathode Rays must be negatively charged!

• 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.

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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 13

(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.

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 10 8 coulomb/g of e • Thomson named the cathode rays electrons.

• Thomson is considered to be the “discoverer of electrons”.

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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 • 1 st model for the structure of the atom was proposed by Thomson based on the following: • Atoms contain small –ve charged particles (e-s) • 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) 16

Rutherford and the Nuclear Atom

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

• If Thompson’s model was correct then any angles.

deflections.  - particles passing through the foil would be deflected by small • Unexpectedly most of the foil with little or nor • A few however, were  particles passed through the deflected a very large angles 17

Rutherford and the Nuclear Atom

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Rutherford and the Nuclear Atom

• Rutherford’s major conclusions from the  particle scattering experiment 1. The atom is mostly empty space.

2. It contains a very small, dense center called the nucleus.

3.

Nearly all of the atom’s mass is in the nucleus.

4. The nuclear diameter is 1/10,000 to 1/100,000 times less than atom’s radius.

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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.

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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.

– 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”.

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Atomic Number

• The atomic number = # of protons in the nucleus.

– Sometimes given the symbol Z .

– On the periodic chart Z is the uppermost number in each element’s box.

• 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 an element of – The number of electrons in a

neutral

equal to the atomic number.

atom is also 22

Nucleon Number and Isotopes

• Nucleon number (formerly Mass number) the symbol A .

is given • A = # of protons + # of neutrons.

– If Z = proton number and N = neutron number – Then A = Z + N 23

• The Standard Notation used to show mass and proton numbers is: Mass number (= p + n) Charge of particle A Z X C Symbol of the atom (= # of p ) A Z

E for example

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C,

6 48 20

Ca,

197 79

Au

• Can be shortened to this symbolism.

14 N, 63 Cu, 107 Ag, etc.

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.

Isotopes of Hydrogen

Protium

(Hydrogen - 1)

Deuterium

(Hydrogen - 2)

Tritium

(Hydrogen - 3) 25

Mass Number and Isotopes

• The stable oxygen isotopes provide another example.

• 16 O is the most abundant stable O isotope.

• How many protons and neutrons are in 16 O?

8 protons and 8 neutrons • • 17 O is the least abundant stable O isotope.

• How many protons and neutrons are in 17 O?

8 protons and 9 neutrons 18 O is the second most abundant stable O isotope.

• How many protons and neutrons in 18 O?

8 protons and 10 neutrons 26

Mass Spectrometry & Isotopic Abundances • Identifies chemical composition of a compound or sample on the basis of the mass-to-charge ratio electrons of charged particles • A gas sample at low pressure is bombarded with high-energy – 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 27

Mass Spectrometry

• There are four factors which determine the extent of deflection: 1 2 3 4 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.

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Mass Spectrometry

A modern mass spectrometer

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Fig. 5-10a, p. 176

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.

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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 28 Si atom?

– The spectrometer will measure the ratio of the mass of an 28 Si atom to 12 C: » Mass of 28 Si atom = 2.331411

Mass of 12 C atom – From this mass ratio, the

isotopic mass

of the » 2.331411 x 12 amu = 27.97693 amu 28 Si can be found: – The mass of the isotope relative to the mass of the C-12 isotope 31

Table 5-3, p. 178

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 33

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 34

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.

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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 36

Atomic Weight

• Example: Naturally occurring Cu consists of 2 isotopes. – It is 69.1% 63 Cu with a mass of 62.9 amu, – and 30.9% 65 Cu, which has a mass of 64.9 amu.

– Calculate the atomic weight of Cu to one decimal place.

atomic weight atomic weight      63 Cu isotope     63.5

amu for copper     65 Cu isotope   37

Atomic Weight

• Example: The

relative atomic mass

of boron is 10.811 amu. The masses of the two naturally occurring isotopes are 5 10 B and 5 11 B, 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 38

Atomic Weight

10.811

amu  10.811

11.009

 amu    (10.013

amu)          1 0 B isotope   10.013

10.013

x

 1 1 B isotope 11.009

x

11.009

x

 11.009

amu

x

 amu 0.198

 -0.996

x

0 .

199 

x

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Atomic Weight

• Note that because

x

is the multiplier for the 10 B isotope, our solution gives us the fraction of natural B that is 10 B.

• Fraction of 10 B = 0.199 and % abundance of 10 B = 19.9%.

• The multiplier for 11 B is (1-

x

) thus the fraction of 11 B is 1-0.199 = 0.801 and the % abundance of 11 B is 80.1%.

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Questions

1. Calculations: Chemistry 9 th Exercises 28 – 38.

Edition, Chapter 4, 2. 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.

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