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ATOMIC STRUCTURE
ATOMIC STRUCTURE
For the last four years, we learned a very simplified model of the atom. It worked very well to explain the chemical properties of the first 20 elements, alkali metals, earth metals, carbon, the halogens and the Noble gases. However, we are about to venture in to the world of the Transition Metals and other atoms that do not obey the Octet Rule. For this, we need a better model...
What has stayed the same?
What has stayed the same?
The Nucleus: As chemists, we really only focus on the protons, as they attract electrons.
Electron Shells: While it would be more correct to think of these as "Energy Levels", we can still imagine electrons existing in "shells" at different distances from the nucleus.
We still (generally) fill these from the inside-out.
The first energy level (shell) still takes 2 electrons.
The second still takes 8 electrons.
After we put 8 in the third energy level, we start using the fourth energy level. However, things get interesting at #21 (Scandium).
What has changed? What is new?
What has changed? What is new?
Sub-levels (sub-shells): The electrons are arranged into sub-levels or sub-shells.
In the first Energy Level, there is only one sub-level (called "1s").
In the second Energy Level, there are two (called "2s" and "2p").
The third and fourth Energy levels each have three ("3s", "3p", "3d", "4s", "4p" and "4d").
Electron configuration: We use the sub-levels to write electron configurations now.
For example, magnesium used to be 2,8,2. Now it is 1s2 2s2 2p6 3s2
Use this interactive Periodic Table to explore the electron configurations of the first 36 elements.
We are going to use our understanding of atomic structure to:
give electron configurations of atoms
give electron configurations of ions
explain properties of transition metals (and their ions)
explain periodic trends (such as electronegativity and atomic radius)
THREE IMPORTANT RULES/PRINCIPLES
THREE IMPORTANT RULES/PRINCIPLES
We use three rules/principles to determine electron configuration:
Aufbau Rule
Aufbau Rule
There is an order to use when filling the orbitals.
Electrons fill orbitals in a way to minimise the energy of the atom.
Pauli Exclusion Principle
Pauli Exclusion Principle
Electrons have opposite spins to allow two to occupy each sub-orbital.
The 1s orbital can hold only two electrons and, when filled, the electrons have opposite spins.
Hund's Rule
Hund's Rule
Half fill an orbital before pairing the electrons, or starting to fill the next orbital.
Nitrogen with "maximum spin" in the 2p orbital.
Oxygen has one more electron than Nitrogen and as the orbitals are all half filled the electron must pair up.
When used together, we can justify the electron configuration of any atom, especially the "exceptions" in the 4th Period: Cr and Cu:
Chromium: [Ar] 3d54s1
Copper: [Ar] 3d104s1
ION FORMATION
ION FORMATION
Ions form because they are more electrically-stable than their parent atom:
full/empty outermost Energy Level
full orbital
half-filled orbital
It requires less energy to remove the electrons that are farthest from the nucleus. This is usually the electrons in the 4th Energy level. After these have been removed, further electrons may be removed from the 3d orbital, or this orbital may be used for bonding with other elements.
ATOMIC/IONIC RADIUS
ATOMIC/IONIC RADIUS
We have to use our understanding of atomic structure to explain some trends we notice on the Periodic Table of Elements.
Atomic radius depends upon two main things:
Effective Nuclear Charge. This is determined by:
nuclear charge (number of protons)
shielding (due to lower energy levels)
Electric Repulsion. This is determined by:
number of electrons in the same energy level
how full orbitals are (filled and half-filled orbitals take up less space than partially-filled orbitals)
When ion form, electrons are gained or lost. This will affect the radius:
cations are smaller than their parent atom
anions are larger than their patent atom
1ST IONISATION ENERGY
1ST IONISATION ENERGY
Ionisation Energy is the amount of energy required to remove one mole of electrons from one mole of atoms (of one element) in its gaseous state. Ionisation Energies have been used to infer the electron configurations of elements.
Ionisation Energies depend upon three main things:
atomic radius
effective nuclear charge (on the valence electrons)
electric repulsion (between electrons)
Key Learning Objectives:
Define 1st Ionisation Energy
List the factors that affect 1st Ionisation Energy
Describe the general trends in 1st Ionisation Energy (across a period and/or down a group)
Explain trends in 1st Ionisation Energy
Justify unexpectedly high 1st Ionisation Energy values
ELECTRONEGATIVITY
ELECTRONEGATIVITY
We met electronegativity briefly last year. It is a relative measure of how tightly an atom attracts electrons in a chemical bond.
We use the Pauling Scale to represent the strength of this attraction. Fluorine is the most electronegative element with a value of 4.0.
The difference in electronegativity can be used to infer the type of bonding between atoms. The higher the difference, the more ionic a bond will be in nature (transfer of electrons). The closer the electronegativity of the two atoms (to each other) the more covalent a bond will be in nature (sharing of electrons).
Key Learning Objectives:
A definition for electronegativity
The periodic trend for electronegativity
The reasons for these trends (relating to atomic structure)
When there is a bond between two atoms, there is often an electronegativity difference (one atom is more electronegative than the other).
This difference gives us an indication of the type of bond that forms between these atoms:
(a) The distribution of electron density in the HCl molecule is uneven. The electron density is greater around the chlorine nucleus. The small, black dots indicate the location of the hydrogen and chlorine nuclei in the molecule. (b) Symbols δ+ and δ– indicate the polarity of the H–Cl bond.
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