Bonding, Structure, Periodicity with a hint of electron configuration Flashcards
1
Q
what is an electron shell?
A
A group of atomic orbitals with the same principal quantum number, n. Also, known as a main energy level
2
Q
What are shells made up of?
A
Atomic orbitals
3
Q
what is an orbital?
A
an orbital is a region around the nucleus that can hold up to two electrons, with opposite spins
4
Q
how many electrons can be held in shell 1?
A
2
5
Q
how many electrons can be held in shell 2?
A
8
6
Q
how many electrons can be held in shell 3?
A
18
7
Q
how many electrons can be held in shell 4?
A
32
8
Q
what shape are s obitals?
A
spherical
9
Q
what shape are p orbitals?
A
dumbell
10
Q
order of size from smallest to largest
A
electron - orbital - subshell - shell
11
Q
what are the three rules when filling orbitals?
A
- Electrons fill the orbitals in increasing energy
- No more than two electrons with OPPOSITE spins can occupy a single orbital
- Electrons spread out occupying as many orbitals with similar energy as possible, before filling orbitals with another electron
12
Q
what are the subshells? how many lectrons do they hold?
A
s = 2 p = 6 d = 10 f = 14
13
Q
what is the special area when filling subshell?
A
4s fills before 3d
14
Q
what is an ionic bond?
A
The electronstatic attraction between positive and negative ions
15
Q
what is a covalnt bond?
A
A covalent bond is the strong electrostatic attraction between a shared pair of electrons and the nuclei of a bonded atom
16
Q
what is a dative covalent bond?
A
A shared pair of electrons in which the bonded pair has been provided by one of the bonding atoms only.
17
Q
what is metalic bonding?
A
The strong electrostatic attraction between positive metal ions and a sea of delocalised electrons.
18
Q
What is electron pair repulsion theory?
A
Pairs of electrons repel each other as far apart from each other as possible
Lone Pairs repel more than bonding pairs.
19
Q
2 bonding pairs: name, angle, example and draw
A
Linear
180°
CO2
20
Q
3 bonding pairs: name, angle, example and draw
A
Triagonal planar
120°
BF3
21
Q
4 bonding pairs: name, angle, example and draw
A
Tetrahedral
109.5°
CH4
22
Q
6 bonding pairs: name, angle, example and draw
A
Octahedral
90°
SF6
23
Q
3 bonding pairs AND 1 lone pair: name, angle, example and draw
A
Pyramidal
107°
NH3
24
Q
2 bonding pairs AND 2 lone pair: name, angle, example and draw
A
Non-linear
104.5°
H2O
25
What are the special cases for the atomic shapes?
Treat double bonds like single bonds
Charge doesnt change anything
26
what do you say if you get a double bond when describing atomic shapes?
Regions of bonding electrons, instead of bonding pairs
27
How do we explain the atomic shapes/ What do we do to work it out?
1. Count and state the number of bonding pairs and lone pairs
2. Say 'Pairs of electrons repel each other as far apart from each other as possible'
3. If lone pairs present, say 'Lone Pairs repel more than bonding pairs'
4. State name of the shape and its bond angle.
28
What is a lone pair?
An outer shell pair of electrons that is not involved in a covalent bond
29
what is a bonding pair?
A pair of electrons shared between two atoms to make a covalent bond
30
what does it mean when something is electron deficient? What does this happen to?
When an atom that forms a covalent bond has less than 8 electrons in its outer shell of electrons, e.g BF3
31
what does it mean when something has expanded its octet
When an atom that forms a covalent bond has more than 8 electrons in its outer shell of electrons e.g SF6
32
Valency of boron
3
33
valency of sulfur
2, 4 or 6
34
valency of chlorine
1, 3, 5 or 7
35
what are the properties of giant ionic lattices?
melting and boiling point -
High mp and bp, strong electrostatic attarction between ions therfore requireing lots of energy to overcome these bonds
Electrical conductivity -
solid - doenst coduct, ions are fixed and cannot move
molten + aqueous - do conduct, ions are free from the lattice and can move
Solubility in water -
Soluble, ions have charge, and therefore can interact with the dipole on water molecules
36
what are the properties of giant metalic lattices?
Melting and boiling point-
Most are high, strong electrostatic attarction between positive metal ions and the delocalised electrons therfore requireing lots of energy to overcome these bonds
Electrical conductivity -
Solid + Molten are conductive, electrons are delocalised and are free to move
37
what is the structure of ionic compounds?
tThey are in a strucutre called a 'Giant ionic lattice' with regulary arranged alternating positive and negative ions, which is often called a latice
38
what is the structure of metals?
All metals have a giant metallic lattice, which is the regualr arrangement of positive metal ions and a sea of delocalised electrons
39
What are the properties of diamond?
Melting and Boiling Points -
Carbon as diamond has high melting and boiling points as it takes a lot of energy to
break the many strong covalent bonds.
Electrical Conductivity -
Diamond doesn't conduct electricity: the electrons are fixed in bonds and cannot move.
Solubility in Water -
Diamond doesn't dissolve in water: there are no charges that can interact with the
dipoles on water molecules.
40
What are the properties of graphite?
Melting and Boiling Points -
Graphite has high melting and boiling points as it takes a lot of energy to break the
many strong covalent bonds
Electrical Conductivity -
Graphite does conduct electricity: the electrons are delocalised and can move between
layers.
Solubility in Water -
Graphite doesn't dissolve in water: there are no charges that can interact with the dipoles
on water molecules.
41
What are the properties of graphine?
High Melting and Boiling Points -
It takes a lot of energy to break the many strong covalent bonds
Electrical Conductivity -
Graphene does conduct electricity: the electrons are delocalised and can move.
Water Insolubility -
Graphene doesn't dissolve in water: there are no charges that can interact with the
dipoles on water molecules.
Strong -
It is the thinnest and strongest material ever made as you have to break strong covalent
bonds to break the layers.
42
what is the structure of diamond?
Diamond has a giant covalent lattice: each carbon makes four covalent bonds to other
carbon atoms in a tetrahedral shape with bond angles of 109.5.
43
what is the strucutre of graphite?
Graphite has a giant covalent lattice structure. Layers have covalently bonded carbon
atoms with a trigonal planar shape and bond angles of 120º. London forces
hold the layers together. Delocalised electrons can move between the layers.
44
what is the strucutre of graphite?
Graphene is a form of carbon made up of a single layer of hexagonally arranged atoms.
45
what are the properties of simple molecular lattice structures with hydrogen bonds
Melting and Boiling Points -
low melting and boiling point (compared to giant compounds like diamond
and NaCl) as it takes a small amount of energy to break the weak forces (hydrogen
bonds, permanent dipole-dipole, induced dipole-dipole)
Electrical Conductivity -
Water does not conduct electricity
46
What is the strucutre of ice?
Ice has a simple molecular lattice structure with hydrogen bonds: each water molecule
oxygen atom forms two strong covalent bonds to hydrogen and two weaker hydrogen
bonds to hydrogen. Only the hydrogen bonds need to break for melting.
47
what are the properties of simple molecular lattice structures with no hydrogen bonds
Melting and Boiling Points -
Simple molecular compounds like iodine have low melting and boiling points as it takes
a small amount of energy to break the weak London forces.
Electrical Conductivity -
Simple molecular compounds like iodine do not conduct electricity: the electrons are
fixed in bonds and cannot move and there are no ions.
Solubility in Water -
Simple molecular compounds like iodine do not dissolve in water: there are no charges
that can interact with the dipoles on water molecules.
48
What is the strucutre of iodine?
Iodine has a simple molecular lattice structure with London forces: strong covalent
bonds between the atoms in I2 molecules, but only weak forces between molecules
(London forces/induced dipole-dipole interactions)
49
What is a polar bond?
Polar bonds are polar covalent bonds where there is a significant difference in
electronegativity between the two bond atoms.
50
what are dipoles?
Dipoles, delta+ and delta- are the partial charges that are present on each member of a polar
covalent bond.
51
what are polar molecules?
Polar molecules have polar bonds because there is an electronegativity difference
and are not symmetrical so the bond dipoles don’t cancel
52
what are non-polar molecules?
Non-Polar molecules could have polar bonds because there is an electronegativity difference
but are symmetrical so the bond dipoles cancel
53
Three types of dipoles
1. Permanent dipoles exist in polar molecules
2. Temporary dipoles form when there is a temporary rearrangement of electrons
3. Induced dipoles are caused when the electrons in a non-polar molecule are
repelled by dipoles in a neighbouring molecule.
54
How do London forces form?
Use words to describe how Induced Dipole-Dipole Interactions or London forces form:
1. Electron rearrangements lead to temporary dipoles
2. These cause induced dipoles by repulsion of electrons
3. London forces are the attractions between the temporary and induced dipoles
55
what to rememember when drawing hydrogen bonds?
LONE PAIRS!!!!!!!!!!!!!!!
56
What are the two anomalous properties of water caused by hydrogen bonding?
Solid water has a lower density than liquid water -
The hydrogen bonds in ice holds the water molecules further apart in fixed positions with lots of space between them. When the ice melts the water molecules can get closer together, so liquid water is denser.
Water has a higher melting or boiling point than similar compounds -
Water has hydrogen bonds between its molecules whereas compounds such as H2S do not. It takes more energy to break the hydrogen bonds.
57
What are the strengths of each intermolecular forces relative to each other?
Hydrogen bonds > permanent dipole-dipole interactions > London forces
58
Why are the London forces between larger molecules stronger?
More points of contact
59
What with the number of electrons causes London forces to be stronger?
London forces (induced dipole-dipole interactions) are stronger between molecules or atoms with more electrons, because the temporary dipoles are more likely to form.
60
What happens to boiling points down group 8 and 7?
Down Group 8 and Group 7, boiling points increase
More electrons
Stronger London forces (induced dipole-dipole interactions)
More energy needed to break them
61
What happens to alkane boiling points as the number of carbons increases?
Alkane boiling points increase with the number of carbon atoms
Longer chains means more points of contact
Stronger London forces (induced dipole-dipole interactions)
More energy needed to break them
62
why do H2O, NH3 and HF have unusually high boiling points?
They have hydrogen bonds
H2S, PH3 and HCl, etc do not
It takes more energy to break the stronger hydrogen bonds than the weaker London
forces (induced dipole-dipole interactions)
63
what causes a polar bond to form?
There must be an electronegativity difference
| The molecule must NOT be symetrical.
64
what happens if there is an electronegativty difference in a moleucles, but it is symetrical.
Bond dipoles cancel out
65
how are modern periodic tables arranged?
The modern periodic table arranges all the elements in order of their proton number.
66
what is a metaloid?
Elements which display properties in between metals and non-metals
67
what is periodicity?
Periodicity is a repeating pattern of properties shown across each periods of the Periodic Table.
68
How do we describe atomic radius across a period?
The number of electron shells doesnt change, but the number of protons increases, so the electrons are held more strongly and the atom gets smaller
69
what is a period?
a horizontal row of elemetns in the periodic table. Elements show trends in properties across a period
70
What is a group?
A vertical column in the periodic table. Elements in a group hagve similar chemical properties and their atoms have the same number of electrons in their outer shells.
71
Group 2 elements are all:
metals with metalic bonding
reactive towards air, water and acids
reducing agents which get oxidised when they react
72
Group 7 elements are all:
non-metals with covalent bonds
simple structures held together by induced dipole-dipole interactions
reactive towards metals and water
oxidising agents which get reduced when they react
73
what is the definition of first ionisation energy?
The First Ionisation Energy is the energy required to remove one electron from each atom in one mole of gaseous atoms to form a mole of gaseous 1+ ions.
74
what is the definition of second ionisation energy?
The Second Ionisation Energy is the energy required to remove one electron from each ion of one mole of gaseous 1+ ions to form a mole of gaseous 2+ ions.
75
in which way are the electrons lost from an atom?
From the outermost shell first
76
Which is bigger first ionisation energy or second ionisation energy? and why?
second ionisation energy, because same number of protons is attracting fewer electrons more tightly and closer to the nucleus so more energy is needed to remove them.
77
how do we write ionisation equations?
X(g) → X+(g) + e-
78
Why are electrons in outer shells are easier to remove than ones on inner shells?
Distance: the nuclear attraction decreases with distance
Shielding: inner electrons are said to shield outer electrons from the nucleus
79
How can we determine the period from the ionisation energies?
Look for the large jumps in energy.
80
Why does successive ionisation energy increase?
each time there are fewer electrons being attracted by the same number of protons, so the remaining protons are held closer to the nucleus and more tightly.
81
What are the two discontinuitys with ionisation energies?
electron pairing and removing an electron from a higher energy level
82
What is the trend of ionisation energy down a group?
Going down the group, there are more shells
The outer electron is further from the nucleus and more shielded
The attraction of the nucleus for the outer electrons decreases
Less energy is needed to remove the outer electron
(This is despite there being more protons in the nucleus)
ionisation energy decreases
83
What is the trend of ionisation energy across a period?
Across the period, there are the same number of shells and similar shielding
There are more protons in the nucleus
The attraction of the nucleus for the outer electrons increases
More energy is needed to remove the outer electron
Ionisation energy increases
84
Why does Be have a higher ionisation energy than B?
Be 1s2 2s2 = s electron is being removed
B 1s2 2s2 2p1 = p electron is being removed
Boron's outer electron is in a higher energy p orbital
It is shielded by the s electrons and less strongly attracted by the nucleus
Therefore less energy is needed to remove it
so ionisation energy decreases
85
Why does N have a higher ionisation energy than O?
N =1s2 2s2 2p3 = [↑↓] [↑↓] [↑ ][↑ ][↑ ]
O =1s2 2s2 2p4 = [↑↓] [↑↓] [↑↓][↑ ][↑ ]
Oxygen's outer electron is paired, whereas nitrogen's is unpaired
Repulsion means less energy is needed to remove it.
So ionisation energy decreases
86
Why does the mp and bp of metals increase across a period?
Increasing melting point as the metallic bonds get stronger: going from Na+ to Mg2+ to Al3+, we get a higher charge and smaller positive ions attracted to more delocalised electrons.
87
Why does silicon have a very high melting point?
A high melting point as it takes a large amount of energy to break the numerous strong covalent bonds holding the giant diamond lattice together.
88
Why does the mp and bp of metals decrease across a period?
Low boiling points as it doesn't take much energy to break the weak IDD forces between the molecules/atoms. The melting points are consistent with larger molecules having more electrons and therefore stronger IDD forces:
S8 > P4 > Cl2 > Ar
OVERALL:
Fewer e-
Weaker IDD force
Not much energy is required to overcome the london forces
