CHEM1101 Midterm: CHEM1101, Topic 1 Atomic Structure and Bonding - Review
Introductory Bonding
Relate the quantum numbers to number and types of orbitals and recognise the
shapes of different orbitals
Be able to populate the orbitals in the ground and excited state of an atom or ion
Be able to write ground state and excited state electron configuration for atoms and
ions
Write Lewis symbols for atoms and ion
Be able be able to draw a Lewis structure for molecules and ions
Use electronegativity differences to identify non-polar covalent, polar covalent and
ionic bonds
Determine formal charges from a Lewis structure and use identify the most favourable
Lewis structures for a molecule or ion.
Recognise exceptions to the octet rule and draw Lewis structures for central atoms
exceeding the octet rule and for central atoms having an incomplete octet.
Understand the VSEPR model and be able to predict it to describe the 3-D shape of
molecules and ions.
Determine whether a molecule or ion is polar or non-polar based on its geometry and
individual bond dipoles.
Be able to identify the hybrid state of atoms in molecules and ions.
Be able to visualise how orbitals overlap to form sigma and pi bonds.
Key Skills
Developed my Erwin Schrödinger (mathematical treatment to incorporate both the wave and
particle characteristics of an electron)
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Electron desity is given by the square of Schrödinger's wave equation and is denoted by
2
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Solving the equation gives a set of wave functions or orbitals and their corresponding energies
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Each orbital describes spatial distribution of electron density
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Escribes the energy level on which the orbital resides
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n ≥1 (positve inegers)
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Correspond to values in the Bohr model
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Principle quantum number: n
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Defines the shape of the orbital
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Allowed values range from 0 to n-1
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Letters are designated to communicate the different values of l
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Angular Momentum Quantum Number: l
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Describes the three-dimensional orientation of the orbital
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Allowed values are integers ranging from -l to l: −l ≤ml ≤l
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Therefore, there can be 1 s orbital, 3 p orbitals, 5 d orbitals and 7 f orbitals
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Magnetic Quantum number: ml
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Spin describes it magnetic field, which affects it energy
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Only has 2 allowed values, +½ and –½.
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Spin Quantum number
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An orbital is described by a set of 3 quantum numbers
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Quantum numbers
TOPIC REVIEW
04 March 2017
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CHEM1101 Page 1
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Value of l for s orbitals is 0
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Spherical shape
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Radium increases with the value on n
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s orbital
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l = 1
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Shape = 2 lobes with a node between them
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p orbitals
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Value of l = 2
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Four of the five d orbitals have 4 lobes, the other resembles a p orbital with doughnut
around the centre
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dorbital
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In one electron H atoms, orbitals on the same energy level have the same energy - called
degenerate orbitals
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In many electron atoms, there are more electrons so electron-electron repulsion
increases
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Therefore, in atoms with more than 1 electron, not all orbitals on the same energy level
are degenerate
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Orbitals in the same sublevel are still degenerate
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This causes energy levels to begin overlapping (4s is lower in the energy than 3d)
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Degenerate orbitals
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Pauli Exclusion Principle
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No to electrons in the same atom can have exactly the same energy
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Every electron in an atom will have a different set of quantum numbers n, l, ml, and ms.
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Usually, it will be the spin quantum number
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Pauli Exclusion Principle
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The way electrons are distributed in an atom is called its electron configuration
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The most stable organization is the lowest possible energy, called the ground state
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A number denoting the energy level
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A letter denoting the type of orbital
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A superscript denoting the number of electrons in those orbitals
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Each component consists of
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Each box in the diagram represents one orbital
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Half-arrows represent the electrons
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The direction of the arrow represents the relative spin of the electron
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Orbital diagrams:
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This means for orbitals in the same sublevel, there must be one electron in each
orbital before pairing and the electrons have the same spin, as much as possible
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For degenerate orbitals, the lowest energy is attained when the number of electrons
with the same spin is maximised
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Hund's Rule
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To determin ethe electron configurations for elements beyond Ar, it is necessary to
satisfy the (n+ l) Rule whioh takes into account electron interactions
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This ruel requires that orbitals with the lowest value of (n+ l) fill first
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If two or more orbitals have the same (n+ l) value, then of these, the orbital with the
lowest value of n fills first
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Orbital
3d
4s
4p
5s
(n+ l)
5
4
5
5
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Therefore, the orbital filling order here is 4s, 3d, 4p, 5s
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Alfbau Principle: (n+ l) Rule
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Elements in the same group of the periodic table have the ssme number of electrons in
there valence shell (valence electrons)
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The filled inner shell electrons are called core electrons. These include completely filled
d or f sublevels
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We write a shorthand version of an electron configuration using brackets around the a
noble gas symbol and listing only the valence electrons
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Condensed electron configurations
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Fill orbitals in increasing rider of energy
Periodic table
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Electron Configurations
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Document Summary
Relate the quantum numbers to number and types of orbitals and recognise the shapes of different orbitals. Be able to populate the orbitals in the ground and excited state of an atom or ion. Be able to write ground state and excited state electron configuration for atoms and ions. Be able be able to draw a lewis structure for molecules and ions. Use electronegativity differences to identify non-polar covalent, polar covalent and ionic bonds. Determine formal charges from a lewis structure and use identify the most favourable. Recognise exceptions to the octet rule and draw lewis structures for central atoms exceeding the octet rule and for central atoms having an incomplete octet. Understand the vsepr model and be able to predict it to describe the 3-d shape of molecules and ions. Determine whether a molecule or ion is polar or non-polar based on its geometry and individual bond dipoles.
