Module 2.2 - Electrons, Bonding and Structure Flashcards
(89 cards)
1
Q
What is the Principal Quantum Number - n?
A
Indicates the SHELL N.O an electron occupies
2
Q
What does a higher n mean?
A
The larger the quantum number, the FURTHER the shell is from the nucleus and the HIGHER the ENERGY level.
3
Q
What energy level and how many electrons does n = 1 have?
A
> 1st Shell
> 2 electrons
4
Q
What energy level and how many electrons does n = 2 have?
A
> 2nd shell
> 8 electrons
5
Q
What energy level and how many electrons does n = 3 have?
A
> 3rd shell
> 18 electrons
6
Q
What energy level and how many electrons does n = 4 have?
A
> 4th shell
> 32 electrons
7
Q
What is an Orbital?
A
Region where electrons may be found
8
Q
Where are orbitals found?
A
They make up energy levels/shells
9
Q
What must the electrons in an orbital be like?
A
> Must be 2 electrons
> Opposite spins
10
Q
What are the 4 types of orbitals?
A
s, p, d, f
11
Q
Describe s-orbitals
A
> From n = 1 upwards, each shell has ONE s-orbital
1 s-orbital with 2 electrons = 2s electrons (max)
Circular shape
12
Q
Describe p-orbitals
A
> From n = 2 upwards, each shell contains THREE P-orbitals:
- Px
- Py
- Pz
3 p-orbitals with 2 electrons = 6p electrons (max)
Three stretched infinity symbols on the x, y and z axis
13
Q
Describe d-orbitals
A
> From n = 3 upwards, each shell contains 5 d-orbitals
5 d-orbitals with 2 electrons in each = 10d electrons
Complex structure
14
Q
Describe f-orbitals
A
> From n = 4 upwards, each shell contains 7 f-orbitals
7 f-orbitals with 2 in each = 14f electrons
Complex structure
15
Q
Electron configuration of Argon - 18
A
1s2 2s2 2p6 3s2 3p6
16
Q
Electron configuration of Zinc - 30
A
1s2 2s2 2p6 3s2 3p6 4s2 3d10
[Ar] 4s2 3d10
17
Q
Why don’t electrons in the same orbital repel?
A
Due to OPPOSITE SPINS.
18
Q
What three principles/rules are important to filling electron orbitals?
A
> Aufbau principle
Pauli Exclusion principle
Hund’s rule
19
Q
What is the Aufbau principle?
A
Electrons must fill from the LOWEST energy
level
20
Q
What is the Pauli Exclusion principle?
A
Electrons (in the same orbital) must have OPPOSITE SPINS
21
Q
What is HUND’s rule?
A
Fill EMPTY orbitals before filling second electrons
22
Q
What is the order of increasing energy level (1s to 4f)?
A
- 1s
- 2s
- 2p
- 3s
- 3p
- 4s
- 3d
- 4p
- 4d
- 4f
23
Q
What is ionic bonding?
A
The electrostatic attraction between positive and negative ions
24
Q
What is the positive ion in ionic bonding?
A
A metal ion is positive
25
What is the negative ion in ionic bonding?
The non-metal ion is negative
26
What happens to the electrons in ionic bonding?
Electrons are transferred from the METAL atom to the NON-METAL atom
27
Explain the solid structures of giant ionic lattices such as NaCl
> Each ion is surrounded by a oppositely charged ion.
> Therefore ions attract each other from all directions.
> Forming 3d giant ionic lattice.
28
Explain the high melting/boiling points of ionic compounds
A large amount of energy is required to break the strong electrostatic bonds that hold oppositely charged ions
29
Why does MgO have a higher melting point than NaCl?
This because the ions in MgO (Mg^2+ and O^2-) have a greater charge than NaCl (Na^+ and Cl^-)
30
Explain why solid ionic compounds cannot conduct electricity
This is because there are only fixed ions which cannot move and therefore cannot carry a current.
31
Explain why ionic compounds conduct electricity when molten or dissolved
This is because the solid lattice is broken down and the fixed ions are now free to move, therefore able to carry a charge.
32
Describe the solubility of ionic compounds
Ionic compounds only dissolve in POLAR solvents such as water. The polar molecules surround each ion to form a solution.
33
What molecules of water surround the positive (metal) ion in an ionic compound?
The negative oxygen molecules surround the positive metal ions.
34
What molecules of water surround the negative (non-metal) ion in an ionic compound?
The positive hydrogen molecules surround the negative non-metal ions.
35
What is covalent bonding?
The strong electrostatic attraction between a shared pair of electrons and the nuclei of bonded atoms
36
How are two positive nuclei able to overcome their repulsion in covalent bonding?
The attraction of the two nuclei to the shared PAIR of electrons overcomes the repulsion
37
Define average bond enthalpy. What does a larger value of bond enthalpy mean?
> Measurement of covalent bond STRENGTH.
| > A stronger covalent bond.
38
What is a lone pair?
An outer-shell pair of electrons that is NOT involved in chemical bonding.
39
What is a dative covalent bond?
When one of the atoms supplies BOTH the shared electrons to the covalent bond.
40
Describe the formation of the ammonium ion, NH4^+
NH3 (lone pair) + H^+ —> NH4^+ (with dative covalent bond at lone pair with H)
41
What is the octet rule?
Eight electrons in the outer shell
42
What happens when there are not enough electrons for an octet?
Unpaired electrons still pair up
43
How is BF3 formed despite having insuffiecient electrons for an octet?`
Three covalent bonds formed with boron three outer electrons which means that six electrons are now in its outer shell, so it doesn’t achieve an octet . The fluorine gains an electron in its outer shell so it achieves an octet.
44
What happens when there are too many electrons to achieve an octet?
More than 4 pairs of electrons end up bonding, so one of the bonding atoms may end up with more than 8 electrons in its outer shell.
45
What is the improved octet rule?
> Unpaired electrons pair up
| > The maximum number of electrons that can pair up is equivalent to the number of electrons in the outer shell
46
What are simple molecular structures made up of?
Made up of small, simple molecules such as H2, O2, Ne, etc
47
Describe solid simple molecular lattice
Atoms in molecules are held by covalent bonds. Different molecules are held together by weak, intermolecular forces (LONDON forces)
48
Describe the electrical conductivity of simple molecular structures
Do not conduct. As there are no free to move ions (fixed)
49
Describe the low melting/boiling points of simple molecular structures
This is because only weak intermolecular forces need to be broken (not strong covalent bonds between atoms IN molecules) so less energy is required.
50
Describe the solubility of simple molecular structures
Soluble in non-polar solvents such as hexane. This is because weak London forces are able to form between the covalent molecules and these solvents. This helps the molecular lattice break down and the substance dissolves.
51
What are giant covalent structures made up of?
Made up of atoms all with covalent bonds.
52
What are some examples of giant covalent molecules?
Diamond, graphite and SiO2
53
Describe the electrical conductivity of giant covalent structures.
Non-conductors as there are no free electrons except graphite which has delocalised electrons
54
Describe the high melting/boiling points of giant covalent structures
This is because all molecules are covalently bonded and all intermolecular bonds must be broken therefore a lot of energy is required.
55
Describe the solubility of giant covalent structures
Insoluble in polar and non-polar solvents as intermolecular bonds are too strong to be broken by polar or non-polar solvents
56
Give: > N.o of areas of electron density
> Bonded pair : Lone pair
> Bond Angle
> Example
of: a linear molecule
> 1 or 2
> 1:0 or 2:0
> 180˚
> H2 or CO2
57
Give: > N.o of areas of electron density
> Bonded pair : Lone pair
> Bond Angle
> Example
of: a trigonal planar molecule
> 3
> 3:0
> 120˚
> BCl3
58
Give: > N.o of areas of electron density
> Bonded pair : Lone pair
> Bond Angle
> Example
of: a tetrahedral molecule
> 4
> 4:0
> 109.5˚
> CH4
59
Give: > N.o of areas of electron density
> Bonded pair : Lone pair
> Bond Angle
> Example
of: a pyramidal molecule
> 4
> 3:1
> 107˚
> NH3
60
Give: > N.o of areas of electron density
> Bonded pair : Lone pair
> Bond Angle
> Example
of: a non-linear molecule
> 4
> 2:2
> 104.5˚
> H2O
61
Give: > N.o of areas of electron density
> Bonded pair : Lone pair
> Bond Angle
> Example
of: a trigonal bipyramid molecule
> 5
> 5:0
> 120˚ and 90˚
> PCl5
62
Give: > N.o of areas of electron density
> Bonded pair : Lone pair
> Bond Angle
> Example
of: a octahedral molecule
> 6
> 6:0
> 90˚
> SF6
63
Give: > N.o of areas of electron density
> Bonded pair : Lone pair
> Bond Angle
> Example
of: a square planar molecule
> 6
> 4:2
> 90˚
> XeF4
64
What is the shape of an ion or molecule determined by?
The number of electron pairs (bonding or lone)
65
Why are the shapes of ions/molecules the way they are?
Allows electrons to be as far apart as possible
66
What is the descending order of strengths of repulsion?
1. Lone/Lone
2. Lone/Bonding
3. Bonding/Bonding
67
Define electronegativity
Ability of an atom to attract a bonded atom for the pair of electrons, in a covalent bond.
68
What happens to electronegativity as group n.o increases?
Increases
69
What happens to electronegativity as period n.o increases?
Decreases
70
What type of dipole do polar bonds have?
Permanent
71
How is a dipole created in a polar bond?
When one bonding atom is more electronegative than the other
72
What is a permanent dipole?
A small charge difference across that results from a difference in electronegativities in bonded atoms.
73
How can a bond be non-polar?
When they are equally electronegative such as H—H
74
How can molecules be polar?
Non-symmetrical so there is a CHARGE DIFFERENCE
75
How can molecules be non-polar?
Symmetrical so charges CANCEL out.
76
What causes intermolecular forces?
Constant random movements of the electrons in shells of atoms of molecules. Don’t involve sharing of electrons
77
What are the two main types of IMF?
> Hydrogen bonding
| > Van der Waals’ forces
78
What are the two main types of Van der Waals’ forces?
> Dipole-dipole interactions (permanent-permanent or permanent-induced)
> London dispersion forces (induced-induced)
79
Describe a permanent-induced dipole interaction
> Caused by a polar molecule being near a non-polar molecule
> Causes electrons to shift in the non-polar molecule
> Therefore making the non-polar molecule slightly polar
80
Describe permanent-permanent dipole interaction
Opposite ends of permanent dipoles are attracted to each other.
H—Cl - - - - - - - H—Cl
81
How are London dispersion forces created?
> Constant random movements of electrons unbalances the distribution of charge
> Causes a instantaneous dipole
> This INDUCES a dipole in neighbouring molecules
> The small induced dipoles then attract causing weak IMF known as London forces
82
What causes a larger London force?
The greater the number of ELECTRONS, the larger the INDUCED dipoles and the greater the attractive forces
83
How is boiling point affected by London forces?
> Substances with London force are low due them being weak
| > The greater the London force, the greater the boiling point.
84
```
Order in terms of relative strength:
> Permanent dipole interactions
> Hydrogen bonds
> London dispersion forces
> Ionic/Covalent bonds
```
1. Ionic/Covalent bonds - Strongest
2. Hydrogen bonds
3. Permanent dipole interaction
4. London dispersion forces - Weakest
85
What is hydrogen bonding?
A strong permanent-permanent dipole attraction between:
> an electron-deficient hydrogen atom (O—H^∂+)
> a lone pair on a highly electronegative atom (O, N or F) on a different molecule
86
What are 4 anamolous properties of water?
> Ice is less dense than water
> Water has higher than expected melting/boiling points
> High surface tension
> High viscosity
87
Why is ice less dense than water?
> Ice has an open lattice with hydrogen holding the water molecules apart
> Water don’t have rigid hydrogen bonds, allowing the water molecules to move closer together.
88
Why is water thought to have an unusually high melting/boiling point?
Other group 16 hydrides have lower melting/boiling points when compared to water
89
Why does water have higher melting/boiling points than other group 16 hydrides?
Hydrogen bonding in water is stronger than IMFs in other group 16 hydrides. Therefore more energy is required to melt/boil water.
