Unit 1 - Chemical Changes and Structure Flashcards
1
Q
How do metallic bonds form?
A
Between metals
Positive nuclei attracted to negative electrons
2
Q
Why do metals conduct electricity?
A
Delocalised electrons
3
Q
What increases the strength of a metallic bond?
A
Increased number of delocalised electrons, increased charge on metal ions, increased strength of metallic bonds
4
Q
What is the structure of metallic bonds?
A
Metal lattice, sea of delocalised electrons
5
Q
Give example of metals that form metallic bonds.
A
Period 1 and 2 (metals)
6
Q
What are monatomic gases?
A
Gases that exist as separate or discrete atoms
7
Q
Give examples of monatomic gases.
A
Helium, Hydrogen, Neon, Argon, Krypton, etc
8
Q
What is a covalent molecular gas?
A
Diatomic elements
9
Q
How do diatomic elements bond?
A
Covalently
10
Q
What is a covalent bond?
A
A shared pair of electrons
11
Q
Give examples of elements that are covalent molecular gases.
A
Diatomic gases
12
Q
What increases the strength of intermolecular bonds?
A
As the size of the atom increases, the number of electrons also increases, increasing the strength of the intermolecular forces between atoms
This increases the melting and boiling points
13
Q
Give examples of covalent molecular solids.
A
Phosphorous (P4)
Sulfur (S8)
14
Q
What is a covalent molecular solid?
A
An element of a set number (P4) bonding with itself covalently
15
Q
What type of intermolecular force do covalent molecular solids have?
A
Weak London Dispersion Forces
16
Q
What increases the strength of an intermolecular force?
A
Increased size = increased number of electrons = increased strength
17
Q
What is a fullerene?
A
Discrete molecules containing 60 or more atoms covalently bonded
18
Q
What does discrete mean?
A
An compound that has a set amount of atoms
19
Q
What is a covalent network solid?
A
Atoms of an element that bond to other atoms of the same element
20
Q
Give examples of covalent network solids?
A
Boron, Carbon (graphite and diamond) and Silicon
21
Q
Why do covalent network elements have high melting and boiling points?
A
Have extremely strong intramolecular covalent bonds that require a lot of energy to break
22
Q
Why do monatomic elements exist as a gas at room temperature?
A
Held together by weak intermolecular forces
23
Q
Why do noble gases have no intramolecular bonds?
A
Have full electron shells so there are no bonding electrons available to form a covalent bond
24
Q
Why do diamonds not conduct electricity?
A
Every carbon atom is bonded to four other carbons and therefore has no delocalised electrons for electricity to pass through
25
What structure do diamonds have?
Tetrahedral
26
What structure does graphite have?
Hexagonal plate arrangement
27
What are the three trends observed in the periodic table?
Covalent radius
Electronegativity
Ionisation Energy
28
What is covalent radius?
Half of the distance between two nuclei of covalently bonded atoms of the same element
29
What is the trend for covalent radius across a period?
Across a period, the number of protons increase, increasing the nuclear charge, attracting the outer electrons closer to the nucleus, decreasing the covalent radius
30
What is the trend for covalent radius down a group?
Down a group, the number of filled electron shells increases, increasing the shielding effect, preventing the outer electrons from being strongly attracted to the nucleus, increasing the covalent radius
31
Why is covalent radius a periodic property?
Follows a clear trend across a period and down a group
32
Why is there no value for covalent radius for the noble gases?
Monatomic so don't bond so the covalent radius can't be calculated
33
What is electronegativity?
A measure of an atom's nuclear attraction for the electrons in a covalent bond
34
What is the trend for electronegativity across a period?
Across a period, the number of protons increase, increasing the nuclear charge, drawing the outer electrons closer to the nucleus, increasing the electronegativity
35
What is the trend for electronegativity down a group?
Down a group, the number of filled electron shells increases, increasing the shielding effect, preventing the outer electrons from being strongly attracted to the nucleus, decreasing the electronegativity
36
Why is electronegativity classed as a periodic property?
Follows a clear trend across a period and down a group
37
Why aren't there electronegativity values for the noble gases?
Monatomic, have full electron shells and so there are no bonding electrons
38
What is ionisation energy?
The energy required to remove an electron from an atom in a gaseous state
39
What is the trend for ionisation energy across a period?
Across a period, the number of protons increases, increasing the nuclear charge, attracting the outer electrons closer to the nucleus, thus increasing the ionisation energy
40
What is the trend for ionisation energy down a group?
Down a group, the number of filled electron shells increases, increasing the shielding effect and preventing the outer electrons from being strongly attracted to the nucleus, decreasing the ionisation energy
41
Is ionisation an endothermic process or an exothermic process?
Endothermic as it requires energy
42
Why is ionisation energy a periodic property?
Follows a clear trend across a period and down a group
43
How do you write ionisation energy equations?
1st IE = Mg(g) --> Mg+(g) + e- Enthalpy change = +738kJmol
2nd IE = Mg+(g) --> Mg2+(g) + e- = +1451kJmol
= +2189kJmol
44
Why is there a large difference in ionisation energies between the first and second ionisations of Lithium?
Second electrons requires the breaking of a full electron shell that is closer to the nucleus
45
Why isn't there a fourth ionisation energy for Lithium?
Lithium only has three electrons
46
What are ionic compounds?
Compounds that form when a metal and a non metal have a large electronegativity difference.
47
What state are ionic compounds in at room temperature?
Solid, high melting point
48
What is the order of states of matter?
Solid -> liquid -> gas
49
In an ionic compound, what loses an electron and what gains the electron?
`Metal loses electron, non metal gains electron
50
When do ionic compounds conduct electricity?
When molten or in solution
51
What state are discrete covalent molecular compounds in?
Liquid or gas
52
In summary, what are the properties of covalent network solids?
High melting and boiling points
| Don't conduct electricity (except graphite)
53
In summary, what are the properties of covalent molecular solids?
Low melting and boiling points
Insulators/don't conduct electricity
P and S
54
In summary, what are the properties of covalent molecular liquids and gases?
Low melting and boiling points
| Insulators/ don't conduct
55
What is a pure covalent bond?
When there is an equal attraction in each atom (same electronegativity values)
56
Where do electrons sit in a pure covalent bond?
In the centre
57
What kind of bond do diatomic elements have?
Pure covalent
| Temporary dipoles
58
What is a dipole?
When one end of a mole becomes slightly positive and the other slightly negative as a result of electrons 'wobbling'
59
How do you compare electronegativity?
The nucleus of __1__ has a greater attraction for bonding electrons than __2__. Electrons will always be nearer __1__. This results in a polar covalent bond because each end has a different charge. The molecule has a permanent dipole
60
What does it mean by 'least polar'?
The compound with the least difference in electronegativity values
61
What is a permanent dipole?
An unequal sharing of electrons between molecules
62
What is the intramolecular bonding continuum?
```
The differences in electronegativity
0 = pure polar covalent
0 - 0.4 = non polar covalent
0.4 - 1.7 = polar covalent
1.7 -> = ionic
```
63
What determines the overall polarity of a molecule?
The symmetry of the charge distribution
64
Are hydrocarbons polar or non-polar molecules?
Non-polar
65
How do you decide if a molecule is overall polar or non-polar?
Draw out individual compounds
Find electronegativity differences
Decide whether polar or non-polar
See if they can cancel out
66
How do you cancel out polar charges when dealing with symmetry?
Take note of polar and non polar bonds
If polar bonds are divisible by 2, they cancel out, making it non polar
(e.g. 3x non-polar bonds, 1x polar bond - don't cancel out - Polar molecule)
(e.g. 2x non-polar bonds, 2x polar bonds - cancel out - Non polar molecule)
(e.g. 4x polar bonds - cancel out - Non polar)
67
Is water polar or non-polar?
Polar
68
Give an experiment used to test whether a liquid is polar or non-polar.
Fill burette with liquid on top of a beaker
Place a charged rod near the stream of water
If the water deflects towards the rod, the liquid is polar.
If the water has a normal path, the liquid is non-polar
69
What are the three types of intermolecular forces, from weakest to strongest?
London Dispersion Forces
Permanent dipole - permanent dipole interactions
Hydrogen bonding
70
What is the alternative name for intermolecular forces?
Van der Waals
71
What is the alternative name for London dispersion forces?
Temporary dipole - temporary dipole interactions
72
What determines the type of intermolecular force present in a compound?
The intramolecular bonds (the polarity)
73
What compounds have LDF's?
Molecules with pure covalent bonds
Monatomic gases
Diatomic elements
Covalent molecular elements
74
What increases the strength of an LDF? Why?
An increase in the number of electrons increases the strength of an LDF as it contributes to the partial charge when dipoles appear
75
What can confirm the strength of an intermolecular force?
The melting and boiling points of noble gases
| low, weak bonds
76
Where do intermolecular forces occur?
Between molecules, between the slightly positive atom of one molecule and the slightly negative atom of another molecule
77
Moving down the halogen group, why does the state change from gases to liquids to solids?
Down the group, the number of electrons increase, creating a stronger partial charge when dipoles occur. This increases the strength of the LDF's and so the molecules are brought closer together, increasing the melting and boiling points
78
Which element has a higher melting and boiling point : Sulphur or Phosphorous? Why?
Sulphur
Sulphur has 128 electrons (S8) while phosphorous have 60 (P4). Sulphur therefor has a stronger partial charge and stronger LDF's
This brings the molecules closer together, increasing the melting and boiling points
79
When do permanent dipole-permanent dipole interaction forces form?
When atoms have difference electronegativity values
80
What increases the strength of a permanent dipole-permanent dipole interaction?
An increase in the difference of electronegativity values
81
When does hydrogen bonding occur?
When a hydrogen atom is bonded to nitrogen, oxygen, or fluorine
82
*Hydrogen bond is not the bond between hydrogen and NOF, but the intermolecular forces between each of the molecules*
83
What is stronger : Intramolecular bonds or Intermolecular bonds?
Intramolecular
84
Why does water have an unusually high melting and boiling point?
Hydrogen bonds present
85
What increases the strength of an intermolecular bond?
The partial charge on the atoms
86
In terms of bonding, what increases the melting and boiling points?
Stronger inter and intramolecular bonds
87
How would you construct an experiment to investigate the factors affecting the viscosity of alcohols?
Select alcohols of similar mass with differing number of hydrogen bonds
(e.g. propanol, propan-1,2-diol, propan-1,2,3-triol)
Fill up to set volume in a measuring cylinder with a ball bearing at the bottom of each flask.
Turn the flask upside down and measure the time taken for the ball bearing travel the distance of the flask.
The longer it takes = the more viscous the alcohol is = the more hydrogen bonds are present = stronger intermolecular forces
88
What is miscability?
How well something mixes
89
What is solubility?
How well something dissolves
90
What substances are the best solvent for non-polar substances?
Non polar substances
91
What substance is the best solvent for polar substances and ionic compounds?
Polar substances
92
Give examples of polar solvents.
Water
| Short chain alcohols (ethanol)
93
What is oxidisation?
A loss of electrons
| Metal
94
What is reduction?
A gain of electrons
| Non metals
95
How do you write a REDOX equation?
Look at pg12 of data book to find equations
(May need to flip it around)
Scale up to get same number of electrons in both equations
Combine equations, leaving out the electrons
96
What is an oxidising agent?
A substance that promotes the oxidisation of another substance by accepting electrons from it. Therefore, as it is gaining electrons, the oxidising agent is being reduced.
97
What is the reducing agent?
A substance which promotes another substance gain its donated electrons. Therefore, the reducing agent is being oxidised as it has lost electrons
98
Where can you find strong reducing agents?
Top right of pg12 of databook
99
Where can you find strong oxidising agents?
Bottom left of pg12 in databook
100
Give examples of oxidising agents.
Hydrogen peroxide
Potassium Permanganate
Sulphur Dioxide
101
What is an everyday use of hydrogen peroxide?
Bleaching of hair and clothes
Teeth whitening
Antiseptic
102
What is an everyday use of potassium permanganate?
Antiseptic
103
What is an everyday use of sulphur dioxide?
Bleaching
104
What is a redox reaction?
A transfer of electrons between atoms
105
How would you write a REDOX equation from scratch?
Should be given a formula of a reactant and a product
Balance the main element
Balance oxygen by adding water
Balance hydrogen by adding hydrogen ions
Balance the overall charge by adding electrons
106
How would you do a redox titration?
Scale for electrons
Combine equations
Mole ratio
Solve
107
Why would a titration not need an indicator?
Self indicating
