2.0 Atomic Structure
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Candidates should be able to:
(a) identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses
(b) deduce the behaviour of beams of protons, neutrons and electrons in an electric field
(c) describe the distribution of mass and charges within an atom
(d) deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (and charge)
(e) (i) describe the contribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number
(ii) distinguish between isotopes on the basis of different numbers of neutrons present
(f) describe the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals
(g) describe the shapes of s, p and d orbitals (see also Section 13) [knowledge of wave functions is not required]
(h) state the electronic configuration of atoms and ions given the proton number (and charge)
(i) explain the factors influencing the ionisation energies of elements (see the Data Booklet) (see also Section 5) (
j) deduce the electronic configurations of elements from successive ionisation energy data
(k) interpret successive ionisation energy data of an element in terms of the position of that element within the Periodic Table
Taken from Chemistry Singapore-Cambridge General Certificate of Education Advanced Level Higher 2 (2019) Syllabus
According to the Rutherford Model,
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Atoms are electrically neutral.
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Atoms are made up of neutron, protons and electrons.
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The positively charged nucleus contains neutrons and protons.
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Negatively charged electrons orbit nucleus in orbitals. (We will get to that)
Definitions
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Atomic/Proton Number is the number of protons in the nucleus of an atom.
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Nucleon/Mass Number is the number of protons and neutrons in the nucleus of an atom.
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Isotopes are atoms of the same element with the same number of protons but a different number of neutrons.
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Isotopes have the same chemical properties as they have the same electronic configuration.
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Isotopes have different physical properties such as melting point and density as they have different masses.
Representation of Atoms
A nitrogen atom has 7 protons, 7 neutrons, 7 Electrons
A boron ion (3- or gained 3e-) has 5 protons, 6 neutrons, 8 electrons.
Sub-atomic particles in electric field
*Angle of Deflection ∝ |Charge|/Mass (Use with y = kx formula)
2.1 Electron Shells
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Electrons of an atom are arranged in shells.
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The closer the shell, the stronger the electrostatic forces of attraction. *
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Each shell is made of smaller subshells, s, p, d, f.
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The number of subshells is dependent on the principal quantum number of the shell. (2nd shell contains 2 subshells, s and p)
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Electrons in the same subshell have the same energy level.
2.2 Orbitals
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Each subshell contains orbitals which are a region of space around the nucleus of an atom where the probability of locating an electron is the highest.
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s subshell contains a single s orbital.
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p subshell contains 3-p orbitals, px, py, pz.
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d subshell contains 5-d orbitals, dxy, dxz, dyz, dx^2-y^2, dz^2.
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f-subshell contains 7 f orbitals.
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Each orbital holds a maximum of 2 electrons.
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As orbitals go further away from the nucleus, they become bigger in size too.
2.2.1 Orbital Energy Levels
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Energy Level generally increases as it becomes further away from the nucleus.
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However, there are exceptions as we go further, 4s orbital is lower in energy level than 3d orbital.
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However, once a 3d orbital has an electron, it becomes lower in energy than 4s orbital. *
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In a given shell, the energy level increases in the order of s, p, d, f.
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Orbitals of the same subshell that have same energy level are known as degenerate orbitals.
2.3 Electronic Configuration
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Aufbau Principle: The added electron will occupy the orbital with the lowest energy first.
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Pauli Exclusion Principle: Each orbital can hold a maximum of 2 electrons of opposite spins.
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Hund's rule of Multiplicity: A analogy would be a bus; would you sit next to someone or in an empty seat? Likewise, when filling a subshell, electrons will occupy an empty orbital singly first before pairing up.
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Some exceptions are chromium and copper. Chromium has a valence configuration of 3d5 4s1 as a half-filled 3d subshell is more stable. Copper has a valence configuration of 3d10 4s1 as a fully filled 3d subshell is more stable.
Type of Representation
or
2)Noble gas core
[He] 2s2 2p4
3)s, p, d, f orbital
1s2 2s2 2p4
2.4 Ionization Energy
The first ionization energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to one mole of gaseous singly charged cations.
The second ionization energy is the energy required to remove one mole of electrons from one mole of gaseous singly charged cations to one mole of gaseous doubly charged cations.
2.4.1 Factors affecting IE
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Nuclear charge effect: Total positive charges of protons in the nucleus resulting in attraction between negatively charged electrons and positively charged protons. Higher the nuclear charge, stronger the electrostatic forces and hence higher energy to remove a valence electron and hence higher IE.
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Shielding Effect: (Only across periods and not down groups). Shielding occurs when there is repulsion between inner shell electrons weakening overall electrostatic forces of attraction causing lower IE. (We will use this for specific parts only, don't worry).
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Atomic radius: As atomic radius increases valence electrons are further away from the nucleus and thus weaker electrostatic forces of attraction, causing lower IE.
2.4.2 IE Trend Across Period
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As a number of proton increases, nuclear charge increases.
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Shielding effect is neglected as electrons are being added to same principle energy shell as we go across the period.
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Hence effective nuclear charge increases and thus stronger electrostatic forces and hence more energy to remove the outermost valence electron and hence higher IE.
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The general trend showing an increase.
2.4.3 IE Trend Down the Group
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As a number of protons increases, nuclear charge increases.
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As the number of electron shell increases, the outermost electron becomes further away from the nucleus.
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Thus, attraction decreases and hence lesser energy to remove the outermost electron causing lower IE.
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The general trend of decreasing trend of IE down the group.
2.4.4 Successive Ionization Energies
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IE increases with each removal of the electron as with a constant positively charged nucleus, the remaining electrons are held more strongly to the nucleus.
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Sharp increases as electrons are being removed from inner shells (refer to 2.4.3).
2.4.5 Exceptions to Trends
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When it comes to ns2 and ns2np1 (Grp 2 to Grp3) configurations there is a decrease as the Group 3 element's valence electron now experiences increase shielding from the fully filled ns2 orbital causing weaker attractions and hence a decrease in IE.
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When it comes to ns2np3 and ns2np4 (Grp 5 to Grp6) configurations there is a decrease as less energy is required to remove the npx electron in the group 6 element as it experiences inter-electronic repulsion arising from 2 electrons occupying the same orbital of the same energy shell.
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