Topic 2 Flashcards
1
Q
Properties of metals
A
- High melting/boiling points
- Conductors of heat and electricity
- Malleable
- Ductile
- Shiny
- Sonorous
2
Q
What does ductile mean
A
That it can be drawn into a wire
3
Q
Characteristics of metallic bonding
A
- Positive ions (cations)
- Delocalised electrons
- Electrostatic forces of attraction
4
Q
What type of structure do metals have
A
Giant lattice structure
5
Q
Why can metals conduct electricity
A
Delocalised electrons are free to move, through the metal
6
Q
Why does Magnesium have a higher melting point than sodium
A
Magnesium forms 2+ ions, losing 2 electrons whereas Sodium forms 1+ ion losing 1 electrons. There is a bigger difference in charges so the electrostatic forces of attraction are stronger
7
Q
When the question states to draw a dot and cross for an ionic structure what does it mean
A
Square brackets
8
Q
What does a complex ion mean
A
The ion contains more than one atom
9
Q
What are the electrostatic forces of attraction between in an ionic substance
A
Oppositely charged ions attracted to eachother
10
Q
What are the electrostatic forces of attraction between in a metal
A
Between cations and electrons
11
Q
Properties of ionic compounds
A
- High melting temperatures
- Brittleness
- Poor electrical conductivity
- Often soluble in water
12
Q
Why do solid ionic compound not conduct electricity
A
There are no delocalised electrons
and the ions are fixed
13
Q
Why are ionic solids brittle
A
As the layers slide, ions of the same charge are side by side and repel one another, causing the ionic solid to break
14
Q
Why are ionic compounds soluble
A
Due to their polarity, as both the positive and negative ions are attracted to water molecules
15
Q
What do you need to compare when looking at ionic compounds
A
- Charge and size of cations and anions
- Ask yourself ‘is the electrostatic force stronger or weaker’
16
Q
How does the radius of an ion effect how much energy is needed to separate the ions?
A
The smaller the ionic radius the more energy needed to separate the ions
17
Q
What does isoelectronic mean?
A
Same number of electrons
18
Q
How would you compare LiF which needs 1031 Kjmol-1 to separate the ions and RbI which needs 628 Kjmol-1 to separate the ions
A
- Look at the charges of the cations, compare them
- Look at the groups, ionic radii
19
Q
Why does LiF need 1031 Kjmol-1 to separate its ions and RbI needs 628 Kjmol-1 to separate the ions
A
- Li+ and Rb+ both are in group 1 and have a charge of +!
-F- and I- are both in group 7 and have a charge of -1 - Li+ has a smaller ionic radius in comparison to Rb
- F- has a smaller ionic radius in comparison to I-
- More electrostatic forces between the ions in LiF than RbI
- More energy is needed to separate the ions in LiF
20
Q
How is a covalent bond formed
A
It is formed between two atoms when an atomic orbital containing a single electron from one atom overlaps with an atomic orbital also containing a single electron of another atom
21
Q
What are the two type of bonds formed
A
Sigma bonds and pi bonds
22
Q
How are sigma bonds formed
A
By end on overlaps
23
Q
How are pi bonds formed
A
By side on overlaps
24
Q
In which atoms can pi bonds exist in
A
only atoms which are joined by double or triple bonds
25
Under what circumstances can a pi bond be formed
It can only be formed when a sigma bond has been formed firstly
26
Is a sigma or pi bond stronger
A sigma bond is a lot stronger
27
How does the weakness of a pi bond relate to alkenes
It is the reason for the increased reactivity of the alkenes and why they can readily undergo addition reactions
28
What is the general trend of bond length and bond strength
The shorter the bonds the stronger the bond strength, this is due to the increase in electrostatic attraction between the two nuclei and the electrons in the overlapping atomic orbitals
29
How many sigma and pi bonds does ethene make
5 sigma bonds
1 pi bond
30
What is the definition of electronegativity
The ability of an atom to attract a bonding pair of electrons
31
What happens to electronegativity along a group
It decreases down a group
32
What happens to electronegativity across a period
It increases from left to right across a period
33
How is the electron density distributed if two atoms of the same element are bonded together by overlapping orbitals
The distribution of the electron density between the 2 nuclei is symmetrical
34
If there is 2 different elements overlapping what type of bond does it form
A polar covalent bond
35
What does a polar covalent bond mean
One side has a slightly positive charge the other side has a slightly negative charge.
It is a covalent bond that has some degree of ionic character
36
How are the contour lines drawn if the electronegativity is higher on one element compared to another
The contour lines are more closely spaced together
37
What is a discrete molecule
It is an electrically neutral group of two or more atoms held together by chemical bonds
38
What is the octet rule
Refers to the element having a full outer shell to be stable, as usually this number is 8
39
How is a dative covalent bond formed
When one empty orbital of one atom overlaps with an orbital containing a non-bonding pair (lone pair) of electrons of another atom
40
How is a dative covalent bond usually drawn and represented
By an arrow going from the atom providing the lone pair of electrons into the empty orbital
41
What does the electron pair repulsion theory state
- the shape of a molecule is caused by the repulsion between the pairs of electrons, both bonded and lone pairs
- lone pair, lone pair repulsion is the strongest
- the electrons arrange themselves to minimise repulsion, maximise separation
42
What is the bond angle for a linear shape
180 degrees
43
What is the bond angle for tetrahedral
109.5 degrees
44
What is the bond angle for v-shaped
104.5 degrees
45
What is the bond angle for trigonal planar
120 degrees
46
What is the bond angle for trigonal pyramidal
107 degrees
47
Examples of a linear molecule
BeCl2, CO2
48
Examples of trigonal planar
BCl3
49
Examples of tetrahedral
CH4
50
Examples of trigonal pyramidal
NH3
51
Examples of V-shaped
H2O
52
What is a dipole
It is said to exist when two charges of equal magnitude but opposite signs are separated by a small distance
53
In relation to dipole what does it mean to be non-polar
Cancellation of dipole charges, overall the molecule will have no dipole
54
In relation to dipole what does it mean to be polar
If the dipole reinforces one another than the molecule will possess an overall dipole, and will be polar
55
Electronegativity is represented by an arrow going towards the higher valued element, what does it mean if the arrows go in opposite directions
The charges cancel each other out and therefore the molecule is non polar
56
Linear, non polar molecule
CO2
57
Trigonal planar, non polar molecule
BCl3
58
Tetrahedral, non polar molecule
CCl4
59
Tetrahedral, polar molecule
CHCl3
60
Trigonal pyramidal, polar molecule
NH3
61
V-shaped, polar molecule
H2O
62
V-shaped and trigonal pyramidal shaped molecules are always, non-polar or polar
POLAR
63
Why do the boiling temperatures increase as you go down the group from chlorine to iodine
- Iodine has more electrons
- so it has greater London forces
- therefore more energy is required to overcome the forces
64
Why are the bond lengths between O-H and S-H different
- Sulfur has a larger atomic radii than oxygen
- therefore sulfur has greater electron shielding
- this reduces the attraction for the bonding electrons
65
Why is the melting temp of sodium sulfide higher than sodium chloride even though both contain ionic bonding
- Sodium sulfate creates a -2 ion
- Sodium chloride creates a -1 ion
- the greater difference in charge creates a larger force of attraction meaning more energy is required
66
What is meant by the term ionic bond
Electrostatic attraction between oppositely charged ions
67
3 types of intermolecular forces
- London forces
- Permanent dipole
- Hydrogen bonds
68
Are intermolecular forces stronger or weaker than covalent and polar covalent bonds
Weaker
69
What causes London forces
An uneven distribution of electrons around a molecule, this causes an instantaneous dipole, one end is negative one is positive, this causes an induced dipole, the two dipoles interact favourable with eachother and is responsible for the london forces of attraction between the molecules
70
What factors increase the overall london forces of attraction
- More electrons as this increases the fluctuation in electron density
- More points of contact between the molecules
71
What is the origin of permanent dipoles
permanent dipole - dipole interaction
72
What is the nature of a hydrogen bond
- The atom bonded to hydrogen has to be more electronegative than hydrogen
- there must be some bond evidence between hydrogen and another atom
73
What are the common elements that experience hydrogen bonds
Oxygen, O
Nitrogen, N
Fluorine, F
74
Important example of hydrogen bonding
Water, H2O
75
What effects the strength of the hydrogen bond
The strength increases as the electronegativity of the element which hydrogen is attached to increases
76
Why can water make two hydrogen bonds
As it has 2 pairs of lone electrons
77
Why does the boiling temp from methane to propane increase
Because the number of electrons increases, so there are stronger London forces
78
Why does hydrogen bonding cause ice to be less dense than liquid water
- More open space between molecules
- due to ring structure in ice
- the hydrogen bonds are longer than the covalent bonds
79
What bond does NH4+ make with one of the hydrogens
A dative covalent bond
80
Why do the boiling temps of unbranched alkanes increase through the first 10
- as the molecular mass increases so does the number of electrons per molecule so the instantaneous and induced dipole increases
- as the length of the carbon chain increases, the number of points of contact between molecules increases, so therefore greater london forces of attraction
81
Why do branched alkanes have a lower melting point than unbranched alkanes
Because the branching means they do not pack well together and have less points of contact and decrease in overall intermolecular forces
82
How does the O-H bond of alcohol affect the boiling point
It forms a hydrogen bond, and it has London forces meaning the forces of attraction is stronger
83
Whats happens to the strength of the hydrogen bond as the chain length of alcohol increase
The strength decreases as the chain length increases
84
Anomalous properties of water
1. Has a relatively high melting & boiling point for a molecule with so few electrons
2. The density of ice at 0'C is less than water at 0'C
85
Why does water have such a high melting and boiling point
The hydrogen bonds between water molecules are relatively strong
86
Why does HF have a lower boiling point than water
HF on average forms one hydrogen bond per molecule and water forms on average 2 so the hydrogen bonding is more extensive
87
Why is the density of ice less than water
- Ice is arranged in rings held together to hydrogen bonds, this creates a large area of open space
- when ice melts the ring structure is destroyed and the average distance between molecules decreases causing an increase in density
88
Why does HF have a higher boiling temp and not fit the general trend of the other hydrogen halides
Due to the unexpected hydrogen bond, which is stronger than intermolecular forces
89
What is the bond angle within a layer of graphene or graphite
All 120 degrees
90
Definition of covalent bonding
Strong electrostatic attraction between two nuclei, with a shared pair of electrons
91
What type of bond joins 2 AlCl3 molecules together
Dative covalent bonds
92
Definition of electronegativity
The ability of an atom to attract the bonding pair of electrons in a covalent bond
93
Explain why 2,2-dimethylpropane has a much lower boiling temperature than its isomer
pentane.
Due to branching which results in weaker london forces due to less points of contact
94
What to look at when comparing ionic bonding
- Differences in charges between ions, is is making a +2 ion
- ionic radii size
95
Why can phosphorous form PCl5 but Nitrogen can only form NCl5
Phosphorous can expand its outershell has it has the 3d orbital available,
Nitrogen does not have a 2d orbital available, it can only fit 8 electrons in its outershell
96
Why does graphite have two different values for compression strength
the lower value relates to the weak london forces between the layers
the higher values relates to the covalent bonding within the layer
