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CHEM1111 > Week 5 > Flashcards

Week 5 Flashcards

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1
Q

Core Electrons

A
  • Electrons in these orbitals are unaffected by the presence of neighbouring atomic nuclei
  • Their energy is practically the same as in an isolated atom
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2
Q

σ or single covalent bonds.

A
  • Electrons in these orbitals are delocalised between neighbouring nuclei.
  • The electron density is highest along the internuclear axis.
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3
Q

Non-bonding (nb) orbitals

A

are localised on only one atom and do not affect bonding.

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4
Q

π bonds

A
  • 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
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5
Q

Total number of atomic orbitals =

A

number of molecular orbitals

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6
Q

Bond order =

A

½ (No. of bonding electrons − No. of antibonding electrons)

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7
Q

From ethane to adamantane

A

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

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8
Q

From adamantine to diamond

A

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

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9
Q

Band gap

A

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.

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10
Q

Silicon

A
  • Small band gap
  • Absorbs all visible light
  • Appears black
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11
Q

Diamond

A
  • Large band gap
  • Absorbs no visible light
  • Appears transparent
  • No conductivity - band gap too big
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12
Q

In order for an electron to conduct electricity…

A

It must have access to an unoccupied energy level.

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13
Q

Allotropes of Carbon

A

Graphite (highly conducting, low density) and diamond (insulator, high density)

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14
Q

Conduction in metals

A

Metals do not have a band gap. Valence and conduction bands overlap, so metals can conduct electricity

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15
Q

Insulator

A
  • Large band gap – electrons cannot be promoted to conduction band
  • Does not conduct electricity, no way the electron could jump
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16
Q

Intrinsic semiconductor

A

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

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17
Q

Extrinsic semiconductor

A

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

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18
Q

n-type doping

A
  • 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
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19
Q

p-type doping

A
  • 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.
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20
Q

Solar cells

A

can be generated by combining an intrinsic semiconductor with both p-type and n-type extrinsic semiconductors

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21
Q

generating solar cells

A
  1. Electrons travel to the conduction band of the n-type semiconductor.
  2. As electrons move, the holes left by electrons in the valence band also change position
  3. The absorption of light promotes electrons from the valence to the conduction band of the intrinsic semiconductor
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22
Q

Spectroscopy

A

The study of the interaction of matter with electromagnetic radiation.

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23
Q

white light is

A

a combination of all colours.
absorb no visible light

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24
Q

black materials

A

absorb all the visible light shone on them

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25
how we see colours due to absorbance
A solid that absorbs red light and reflects all the other colours will appear green
26
Transparent
Materials allow **all** light to pass through
27
Translucent
Materials allow **some** light to pass through, but the light is **scattered**
28
Opaque
Materials **do not** let any light pass through. Light is **reflected** or **absorbed**
29
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
30
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
31
A = εcl
**Beer-Lambert** law A = absorbance. c = concentration. ε = molar extinction coefficient. l = path length.
32
Absorbance
related to how much **light** can pass through a solution. * **Higher** absorbance → **less** light gets through (a darker solution)
33
A more concentrated solution...
will absorb more light
34
Molar extinction coefficient
for a given molecule, this is a **constant**
35
path length
* How **far** the **light** has to travel. * If light has to travel through **more solution**, then **more** will be **absorbed**
36
Atomic absorption spectroscopy (AAS)
uses the characteristic absorption **wavelengths** of each element to *determine* **concentrations** of elements
37
The Beer-Lambert law tells us that A vs c
should be a straight line
38
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
39
elements with **low** ionisation energy
**Metals** * easily *lose* electrons
40
elements with **high** ionisation energy
**Non-metals** * readily *gain* electrons * apart from noble gases
41
each element has an electronegativity on a scale of
0 to 4
42
Atoms *gain* or *lose* electrons to become **isoelectronic** to the nearest **noble gas**
True
43
More than one electron may be gained/lost but >3 electrons are not common
True
44
good electron donors are...
on the **left** of the periodic table * s1, s2 and d (transition elements) * e.g Mg → Mg^2+
45
good electron acceptors are..
on the **right** of the periodic table * e.g. F → F^-
46
ionic radius
an estimate of the size of an **ion** in a ***crystal* lattice**
47
**Cations** are larger than their parent atoms
False, they are **smaller**.
48
**Anions** are smaller than their parent atoms
False, they are **larger**
49
Ionic bonding
**Long**-range **electrostatic** attraction between *cation* (+) and *anion* (−), together with the **short**-range repulsion between **electrons** in *adjacent* ions
50
The **equilibrium** distance between cation and anion nearest-neighbours occurs when...
the potential energy is a **minimum**
51
Electrostatic interactions are isotropic
They are the **same** in all directions – and they are long-ranged (decay as r^−1)
52
How ionic crystals form
1. Shell of oppositely charged ions (**counterions**) is attracted to surround a **central ion**. 2. This in turn attracts **another** shell of their counter ions etc… 3. Structure continues to **grow** leading to **ionic crystals** rather than small molecules.
53
Ionic crystals
an organised **lattice** of **cations** and **anions**. Many different arrangements of ions can form depending on *ionic radii*
54
What makes an ionic crystal stable?
The attraction between **oppositely** charged ions
55
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)
56
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
57
Lattice energy
the energy *change* when **gas phase** ions combine to form a **crystal lattice**
58
Negative lattice energy indicates
that the energy of the crystal lattice is **lower** than that of the ions
59
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
60
Lattice energy depends on a number of factors:
1. The **sum** of cation and anion **radii** 2. The **charge** on the ions 3. The **arrangement** of ions
61
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**
62
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**
63
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**

CHEM1111 (15 decks)

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