Week 5 Flashcards
Core Electrons
- Electrons in these orbitals are unaffected by the presence of neighbouring atomic nuclei
- Their energy is practically the same as in an isolated atom
σ or single covalent bonds.
- Electrons in these orbitals are delocalised between neighbouring nuclei.
- The electron density is highest along the internuclear axis.
Non-bonding (nb) orbitals
are localised on only one atom and do not affect bonding.
π bonds
- Electrons in these orbitals lower the energy of the molecule.
- They are delocalised between multiple nuclei in lobes on opposite sides of the internuclear axis.
- They are responsible for double and triple bonds
Total number of atomic orbitals =
number of molecular orbitals
Bond order =
½ (No. of bonding electrons − No. of antibonding electrons)
From ethane to adamantane
increasing # of atoms, increases # of molecular orbitals.
1 C-C σ anf 1C-C σ star becomes
12 C-C σ and 12 C-C σ star
From adamantine to diamond
Keep growing, so many energy levels it becomes like a continuom.
Refer to the levels as a band
occupied levels = valence band
unoccupied levells = conduction band
Band gap
Gap between the valence and conduction band
* Minimum energy a network solid must absorb to promote an electron from the valence band to the conduction band.
Silicon
- Small band gap
- Absorbs all visible light
- Appears black
Diamond
- Large band gap
- Absorbs no visible light
- Appears transparent
- No conductivity - band gap too big
In order for an electron to conduct electricity…
It must have access to an unoccupied energy level.
Allotropes of Carbon
Graphite (highly conducting, low density) and diamond (insulator, high density)
Conduction in metals
Metals do not have a band gap. Valence and conduction bands overlap, so metals can conduct electricity
Insulator
- Large band gap – electrons cannot be promoted to conduction band
- Does not conduct electricity, no way the electron could jump
Intrinsic semiconductor
Band gap is small, electrons can be promoted to the conduction band, leaving electrons in the conduction band and holes in the valence band.
* e.g. silicon, germanium.
* Allows electrons in the valence band to jump in this hole
* Conductivity in both bands
Extrinsic semiconductor
Many applications require stable conductivity at all temperatures.
* This can be achieved by doping – substituting some atoms (something added to the material).
* There are two types of doping – n-type and p-type
n-type doping
- There are extra negative charge carriers, (i.e. electrons).
- Achieved by substituting with an element to the right on the periodic table, which has more electrons.
- Don’t need as much energy
- Extra electrons reside in donor levels, just below the conduction band, heat moves them into this band
p-type doping
- Fewer electrons, and more positive charge carriers, (i.e. holes).
- Achieved by substituting with an element to the left on the periodic table, which has fewer electrons.
- The electron poor atoms generate acceptor levels, just above the valence band, valence band electrons are promoted here, leaving holes in the valence band.
Solar cells
can be generated by combining an intrinsic semiconductor with both p-type and n-type extrinsic semiconductors
generating solar cells
- Electrons travel to the conduction band of the n-type semiconductor.
- As electrons move, the holes left by electrons in the valence band also change position
- The absorption of light promotes electrons from the valence to the conduction band of the intrinsic semiconductor
Spectroscopy
The study of the interaction of matter with electromagnetic radiation.
white light is
a combination of all colours.
absorb no visible light
black materials
absorb all the visible light shone on them
how we see colours due to absorbance
A solid that absorbs red light and reflects all the other colours will appear green
Transparent
Materials allow all light to pass through
Translucent
Materials allow some light to pass through, but the light is scattered
Opaque
Materials do not let any light pass through. Light is reflected or absorbed
Atomic spectrum of hydrogen
- Composed of discrete wavelengths, or “spectral lines”.
- Shows that the gaps between energy levels are fixed, and was early evidence that the energy of the electron in the atom is quantised
Energies of one-electron atoms
- Hydrogen has one electron, so energy levels are defined only by the principal quantum number, n.
- Can be calculated using the Rydberg formula
A = εcl
Beer-Lambert law
A = absorbance.
c = concentration.
ε = molar extinction coefficient.
l = path length.
Absorbance
related to how much light can pass through a solution.
* Higher absorbance → less light gets through (a darker solution)
A more concentrated solution…
will absorb more light
Molar extinction coefficient
for a given molecule, this is a constant
path length
- How far the light has to travel.
- If light has to travel through more solution, then more will be absorbed
Atomic absorption spectroscopy (AAS)
uses the characteristic absorption wavelengths of each element to determine concentrations of elements
The Beer-Lambert law tells us that A vs c
should be a straight line
Molecular spectroscopy
- Does not require atoms to be atomised, as it measures the energy of electrons in molecules, not atoms
- Molecules absorb specific wavelengths of light according to orbital energies.
- Beer-Lambert law applies
elements with low ionisation energy
Metals
* easily lose electrons
elements with high ionisation energy
Non-metals
* readily gain electrons
* apart from noble gases
each element has an electronegativity on a scale of
0 to 4
Atoms gain or lose electrons to become isoelectronic to the nearest noble gas
True
More than one electron may be gained/lost but >3 electrons are not common
True
good electron donors are…
on the left of the periodic table
* s1, s2 and d (transition elements)
* e.g Mg → Mg^2+
good electron acceptors are..
on the right of the periodic table
* e.g. F → F^-
ionic radius
an estimate of the size of an ion in a crystal lattice
Cations are larger than their parent atoms
False, they are smaller.
Anions are smaller than their parent atoms
False, they are larger
Ionic bonding
Long-range electrostatic attraction between cation (+) and anion (−), together with the short-range repulsion between electrons in adjacent ions
The equilibrium distance between cation and anion nearest-neighbours occurs when…
the potential energy is a minimum
Electrostatic interactions are isotropic
They are the same in all directions – and they are long-ranged (decay as r^−1)
How ionic crystals form
- Shell of oppositely charged ions (counterions) is attracted to surround a central ion.
- This in turn attracts another shell of their counter ions etc…
- Structure continues to grow leading to ionic crystals rather than small molecules.
Ionic crystals
an organised lattice of cations and anions. Many different arrangements of ions can form depending on ionic radii
What makes an ionic crystal stable?
The attraction between oppositely charged ions
Packing arrangement of NaCl
Face-centred cubic (fcc) array
* Cl− at corners and faces of the cube; Na+ in spaces in between), due to a large difference in ionic radii.
* Smaller Na+ (1.02 Å) can fit into the spaces (interstices) between closely packed larger Cl− (1.81 Å)
Packing arrangement of CsCl
Primitive cubic array
* Cs+ at centre; 8 Cl− anions at corners, due to similar ionic radii.
* Cs + (1.61 Å) is big enough to fit more than 6 Cl− anions around it.
* Less stable than the fcc array
Lattice energy
the energy change when gas phase ions combine to form a crystal lattice
Negative lattice energy indicates
that the energy of the crystal lattice is lower than that of the ions
The energy of a crystal lattice is due only to…
electrostatic interactions between the ions, so we can calculate it exactly by adding up all of the pairs of interactions between ions as long as we know their distances
Lattice energy depends on a number of factors:
- The sum of cation and anion radii
- The charge on the ions
- The arrangement of ions
Once an atom loses an electron to form a cation
The remaining electrons experience a greater attraction to the nucleus due to the charge imbalance and so the ionic radius decreases
When an atom gains an electron to form an anion
the nucleus exhibits a decreased hold on the electrons owing to the charge imbalance and so the radius increases
When an atom gains an electron to form an anion
the nucleus exhibits a decreased hold on the electrons owing to the charge imbalance and so the radius increases
