Worksheet 2.2
Chapter 2: Atomic structure – fast facts
2.1 The atom
• Protons and neutrons are present in the nucleus of an atom; electrons are in orbits or energy
levels around the nucleus.
•
The relative masses and relative
charges of the sub-atomic
particles are:
Relative mass
Relative charge
Proton
1
+1
Neutron
1
0
Electron
5 × 10
–4
–1
•
Atomic number (Z) = number of protons. It is the fundamental characteristic of an element.
•
Mass number (A) = number of (protons + neutrons).
•
Isotopes are atoms with the same atomic number, different mass number OR the same number
of protons, but different number of neutrons.
•
For a species
:
number of protons = Z
number of electrons = Z – q
number of neutrons = A – Z.
•
Isotopes differ in physical properties that depend on mass such as density, rate of diffusion, etc.
This difference is very significant for the isotopes of hydrogen as deuterium
has twice the
mass of the more abundant
same chemical properties.
•
. As isotopes have the same electron arrangement they have the
Examples of the uses of radioisotopes: C-14 in radiocarbon dating, Co-60 in radiotherapy and
I-131 and I-125 as medical tracers.
2.2 The mass spectrometer
• Stages of operation: vaporization of sample, ionization to produce M + ions by electron
bombardment, acceleration of positive ions by electric field, deflection of ions by magnetic
field perpendicular to their path, detection of ions.
•
Degree of deflection depends on the charge/mass ratio: the smaller the mass or the greater the
charge: the greater the deflection.
•
For an element, the mass spectrum gives two important pieces of information: the number of
isotopes, and the abundance of each isotope. This allows the relative average atomic mass, Ar
to be calculated.
1
•
Relative atomic mass (Ar) of an element is the average mass of an atom according to the
relative abundances of its isotopes, on a scale where the mass of one atom of
is 12
exactly.
For example for Cl which has two isotopes
(75 %) and
(25 %).
•
For a molecule, the peak with largest mass represents the molecular (parent) ion and its mass
gives the relative molecular mass (Mr) of the compound. The fragmentation pattern can help
determine molecular structure (see Chapter 12).
2.3 Electron arrangement
• The electromagnetic spectrum includes waves in order of decreasing frequency/energy, γ rays,
X-rays, ultraviolet radiation, visible light, IR radiation, microwaves, and radio waves. (See
Table 3 of the IB Data booklet).
•
Frequency (f) and wavelength (λ) are related by: c (speed of light) = fλ.
•
The energy of a photon (Ephoton) is related to the frequency (f) of the radiation by Planck’s
equation:
Ephoton = hf (The equation is given in Table 1of the IB Data booklet).
h is Planck’s constant (Table 2 of the IB Data booklet).
•
A continuous spectrum contains light of all wavelengths in the visible range.
•
A line spectrum consists of a few lines of different wavelengths/frequencies.
The lines in an emission spectrum are produced by excited electrons falling from higher to
lower energy levels: ΔEatom = hf = hc/λ.
As the energy levels converge at higher energy as they are further from the nucleus; the lines in
the spectrum also converge at higher energy/frequency.
•
The hydrogen spectrum:
Series
Region
Electron falls to
Lyman
UV
n=1
Balmer
Visible
n=2
IR
n = 3–
Paschen
•
The ionization energy of hydrogen corresponds to the convergence limit of the Lyman series.
2
•
The electron arrangement indicates the number of electrons in each energy level.
Element
Electron
arrangement
Element
Electron
arrangement
H
1
Na
2, 8, 1
He
2
Ar
2, 8, 8
Li
2, 1
K
2, 8, 8, 1
Ne
2, 8
Ca
2, 8, 8. 2
12.1 Electron configuration
• The electron configuration of an atom describes the number of electrons in each energy sublevel.
•
Evidence for the existence of main energy levels and sub-levels comes from graphs of first
ionization energies of successive elements or successive ionization energies of the same
element.
•
The first ionization energy is the minimum energy required to remove one mole of electrons
from a mole of gaseous atoms to form a mole of univalent cations in the gaseous state. It is the
enthalpy change for the reaction: X (g) ↓ X + (g) + e–.
•
Large increases in successive ionisation energies of an atom occur when an electron is removed
from a different energy level. Smaller increases occur when an electron is removed from a different
sub-level. A very small jump occurs when there is a change from a p4 to p3 configuration as paired
electrons are easier to remove than unpaired electrons.
•
The main energy levels (in order of increasing energy) are identified by integers,
n = 1, 2, 3, 4… Each main energy level contains n sub-levels and n2 orbitals.
•
The sub-levels (in order of increasing energy) and orbitals are identified by letters: s, p, d, f,
etc.
•
Orbitals are regions in space in which an
electron may be found in an atom. They have
characteristic shapes. s orbitals are spherical and
p orbitals are dumb-bell shaped. There are
corresponding p orbitals orientated along the x
and z axis.
s orbital
py orbital
•
Pauli Exclusion Principle states that only electrons with opposite spin can occupy the same
orbital.
•
The Electrons-in-boxes notation is used to describe the number of electrons in each orbital.
Each orbital is represented by a box and each electron by an arrow which represents the
direction of its spin.
3
The relative energies of the sub-levels and their composition are summarised.
Level
n=4
n=3
n=2
n=1
Sublevel
Maximum no. of
electrons in sublevel
4f
14 (seven f
orbitals)
4d
10 (five d orbitals)
4p
6 (three p orbitals)
4f
14 (seven f
orbitals)
3d
10 (five d orbitals)
3p
6 (three p orbitals)
3s
2 (one s orbital)
2p
6 (three p orbitals)
2s
2 (one s orbital)
1s
2 (one s orbital)
Maximum
no. of
electrons
in level
32
18
8
2
A useful mnemonic to the order of filling orbitals. Follow the arrows to see the
order in which the sub-levels are filled.
1s, 2s, 2p, 3s, 3p, 4s, 3d. 4p, 5s, 4f, 5d, 6p, 7s…
•
The Aufbau principle states that orbitals with lower energy are filled
before those with higher energy.
•
Hund’s rule states that every orbital in a sub-level is singly occupied
with electrons of the same spin before any one orbital is doubly occupied.
The number of electrons in a sub-level is represented by a super script number.
Element
Electron
configuration
1
H
1s
Li
1s 2s
2
1
B
1s 2s 2p
2
2
1
Na
1s 2s 2p 3s
2
2
6
1
Element
Electron
configuration
1
2
5
1
8
2
Sr
[Ar] 3d 4s
Cr
[Ar] 3d 4s
Ni
[Ar] 3d 4s
Cu
[Ar] 3d 4s
10
1
Note the exceptional configuration of copper and chromium.
4