TOPIC 2: ATOMIC STRUCTURE
2.1
The atom
Atomic Theory before Bohr
Democritus (420 BCE) proposed that matter was made up of small, indivisible particles called atoms
John Dalton (1808)
Billiard-Ball Model
Atoms are the smallest particle of matter → invisible, indivisible & indestructible spheres
Atoms of each element are identical to each other and unique from all other elements
JJ Thompson (1897)
discovered “Cathode Rays” from the cathode (−) plate to the anode (+) plate
The mass of the particles (electrons) were nearly 2000 times smaller than a hydrogen atom
 evidence that there are particles smaller than the smallest atom (H)
Ernest Rutherford (1895)
discovered positively charged α particles and β rays emitted by uranium
α = alpha particle = helium-4 nucleus = 42 He2+
β = beta rays = stream of electrons = e–
The Geiger-Marsden Experiment
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• High energy α particles were directed at extremely thin gold foil surrounded by a fluorescent screen
• Nearly all of the α particles went nearly straight through the gold foil – less than 10° deflection
 nearly all of the gold is empty space
• The gold had a positive charge but no change in mass after the α particle bombardment
 negative charges were removed from the gold by the α particles
 but no α particles were captured by the gold (no change in mass)
• Some of the α particles were deflected more than 90°
 the α particles must have been repelled by a positively charged mass (same charge as the α particle)
2.1.1
State the position of protons, neutrons and electrons in the atom.
Rutherford
developed the Nuclear Model of the atom based on the experiment of Geiger & Marsden
Nucleus =
 the positively charged centre of the atom
 having almost no volume and
 nearly all of the mass
Electrons = the negatively charged particles that circulate around the nucleus
2.1.2
State the relative masses and relative charges of protons, neutrons and electrons.
Unified Atomic Mass Unit
a unit of mass equal to exactly
1
of the mass of one carbon-12 atom.
12
1 atom 126 C = 12 u (exact value)
€ mass, m / kg
1.673 × 10–27
1.675 × 10–27
9.110 × 10–31
p+
proton, €
neutron, n0
electron, e–
2.1.3
mass, m / u
1.007276 ≈ 1
1.008665 ≈ 1
0.000549 ≈ 0
charge
+1
0
–1
Define the terms mass number (A), atomic number (Z) and isotopes of an element.
Mass Number (A)
mass of the nucleus in u = nucleon number
A=p+n
Atomic Number (Z)
nuclear charge = number of protons
Z=p
Neutron number (N) number of neutrons
Isotopes
2.1.4
2.1.5
N=A–Z
atoms of an element having different masses
same number of protons
different number of neutrons
same atomic number (Z)
different mass number (Z)
Deduce the symbol for an isotope given its mass number and atomic number.
Calculate the number of protons, neutrons and electrons in atoms and ions from the mass number, atomic number and charge.
Chemical symbol of an isotope
→ shows the chemical symbol of the element (X)
→ shows atomic mass number (A) as superscript
→ shows atomic number (Z) as subscript
A
ZX
Isotope Name = element name – mass number (carbon-12)
isotope name
hydrogen-1
hydrogen-2 (deuterium)
hydrogen-3 (tritium)
carbon-12
carbon-14
potassium-40
protons
1
1
1
6
6
19
neutrons
0
1
2
6
8
21
symbol
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1
1H
2
1H
3
1H
12
6C
14
6C
40
19 K
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2.1.6
Compare the properties of the isotopes of an element. Discuss the uses of radioisotopes
Physical Properties
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some (related to mass) change slightly
melting point
boiling point
density
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Chemical Properties
stay the same
ratios in compounds (1N:3H → NH3)
reactivity
atomic radius