Electrons, bonding and structure Flashcards
(102 cards)
1
Q
What is the principal quantum number
A
Each shell is given a number e.g closest shell to nucleus is called n=1 and increases by 1 outwards by each shell n=2, n=3 etc
2
Q
What equation is used to find the maximum amount of electrons a shell can hold
A
Maximum amount of electrons a shell can hold = 2n^2
3
Q
What is an atomic orbital
A
It is a region around the nucleus that can hold up to two electrons with opposite spins
4
Q
Key concepts of atomic orbitals:
A
Electrons can either spin up or down- therefore two electrons in the same orbital must oppose each other (i.e have opposite spins)
The negative charge cloud has the shape of the orbital occupied by the electron. However, the exact location of the electron cannot be found (95% probability of where it exists)
5
Q
What is the electron considered to be
A
An electron is considered to be a cloud of negative charge
6
Q
What are the letters of the several types of orbitals
What shells do each orbital appear in
What is the maximum amount of electrons each shell holds
What are the shapes of the orbitals you are required to know
A
1) S,p,d,f
2) S-orbitals appear at the start of every shell
P-Orbitals appear from the second shell after the S-orbital and every other shell after
D Orbitals appear from the third shell after the P-orbital and every other shell after
F Orbitals appear from the fourth shell after the D-orbital and every other shell after
3) 1=2e / 2=8e / 3=18e / 4=32e
4)
S-orbitals are spherical
P-orbitals are shaped as a figure of eight
7
Q
What is a sub shell
A
A sub shell is all of the orbitals of the same type in the same shell
8
Q
Key ideas behind Electron configuration:
A
Different sub shells have different energies. As we move away from the nucleus , the energy of the subshells increases
The energy of the 4s sub shell is less than the energy of the 3d sub shell. Therefore, we must fill the 4s subshell before the 3d subshell. However, the electron configuration is always written in the order of electron shells not filling (4s after 3d)
9
Q
What rules are followed for filling atomic orbitals
A
Orbitals with the lowest energy are filled first
we can have up to two electrons in the same orbitals but they must have opposite spins
If we have orbitals with the same energy, then we put electrons into individuals orbitals before we pair them. That is because electrons in the same orbital repel.
10
Q
Explain how to use shorthand notation
A
Find the element in the periodic table
We then find the noble gas before the element.
The inner electrons shell of the element will be similar to the noble gas.They are represented as the chemical symbol of the noble gas with square brackets surrounding it.
After we would write the rest of the electron configuration. e.g the outer shell involved in the reaction .
11
Q
Electron configuration and shorthand electron configuration of vanadium
A
1s^2 / 2s^2 / 2p^6 / 3s^2 / 3p^6 / 4s^2 / 3d^3
[Ar] 4s^2 / 3d^3
12
Q
How does Chromium and Copper seen to be different in their electron configuration
A
In both cases, the 4s subshell contains only one electron even through there are electrons in the 3d subshell
13
Q
Explain why Chromium and Copper seen to be different in their electron configuration
A
The 3d subshell is more stable when it is either half full or full.
In the case of chromium:
One electron moves from the 4s subshell to form a half full 3d subshell
In the case of copper:
One electron moves from the 4s subshell to form a full 3d subshell
This arrangement increases the overall stability of the atom
14
Q
How is the periodic table divided
A
Arrangement of elements:
-The periodic table is arranged in periods (rows) and groups (columns)
-Elements are ordered by increasing atomic number (number of protons).
Repeating Trends Across a Period:
-Across a period (left to right), elements show repeating trends in their physical and chemical properties.
-This is because the elements in a period have the same number of electron shells but increasing numbers of electrons and protons.
Certain groups:
-On the left side there are the metals
-On the right side are the non-metals
-Between the metal and non-metals are teh metalloids
Arrangement of blocks:
Each block is named after the subshell containing the highest energy electron for the elements in that block
15
Q
How can we use the blocks in the periodic table to determine if the electronic configuration is correct.
A
we count from hydrogen then count across and go down each period while counting the electron in each subshell until you have reached the desired element
16
Q
Explain the complications of the d-block in the periodic table
A
-The first row of the d-block is in period 4. However, the 3d subshell is being filled.
-This is because the 4s subshell fills before the 3d subshell as it is lower in energy Therefore , the 3d subshell fills on period 4
17
Q
Key ideas of shorthand electron configuration
A
The only electron in the outer shell are involved in chemical reactions.(Therefore electrons in the inner shells are not involved in chemical actions. )
Rather than writing down all the electrons scientists use shorthand electron configuration as it is more simple to write.
You will always present full electron configuration in exams unless asked otherwise
18
Q
What happens when an electron is removed
A
An ion is formed with a single positive charge
19
Q
Draw the equation of the formation of a magnesium ion
A
Mg —> Mg 2^+ + 2e^-
20
Q
Define first ionisation energy
A
First ionisation energy is the energy needed to remove one mole of electrons from one mole of atoms in their gaseous state to form one mole of 1+ ions
21
Q
What is shown through this equation in terms of ionisation energy :
Mg(g) —> Mg(g)+ e-
A
One mole of Mg is taken and converted into a gas shown through (g)-gaseous state
One electron from every atom is taken to form one mole of 1+ ions also in their gaseous state
The energy needed to do this is called first ionisation energy
22
Q
What happens after the first ionisation energy
A
Once one electron is removed we can continue to remove electrons and measure the ionisation energy each time.
When another electron is removed the energy required to do this is called second ionisation energy.
23
Q
Define second ionisation energy
A
The second ionisation energy is the energy needed to remove one mole of electrons from one mole of 1+ ions in their gaseous state to form one mole of 2+ ions (also in gaseous state)
24
Q
What is called successive ionisation energies
A
When the electron is continually removed and the ionisation energy is measured each time
25
Differences between first and second ionisation energy
First ionisation energy is the energy required to remove one mole of electrons from one mole of atoms (in their gaseous state) to form one mole of 1+ ions (in their gaseous state)
Second ionisation energy is he energy required to remove one mole of electrons from one mole of 1+ ion (in their gaseous state) to form one mole of 2+ ions (in their gaseous state)
It takes more ionisation energy for the second ionisation energy compared to the first ionisation energy
26
How is it easy to work out which ionisation energy is shown by an equation like this:
Mg(g)4+ —> Mg(g)5+ = e-
The ionisation energy shown is the same as the charge on the ion which will gain an electron
The fifth ionization energy is shown as there is an ion with a 5+ charge which will gain an electron
27
How is ionisation energy determined
The atomic radius: as the atomic radius increases the attraction between the nucleus and outer electrons decreases
Charge on the nucleus: as the number of protons increases the attraction between the nucleus and outer electrons also increases
Shielding: Electrons in the outer shell are repelled by electrons in the inner shells. As the number of inner shells increases the attraction between the nucleus and the outer electrons decreases
28
Explain why the first ionsation decreases when going down a group
The atomic radius increases therefore the outer electron shell is further away from the nucleus
The number of internal energy levels also increases. Therefore, there is more shielding between the nucleus and the outer electrons
Both these factors mean that going down a group the attraction between the nucleus an outer electrons decreases. This causes the first ionisation energy to fall.
Despite the nuclear charge increasing as we go down a group this factor is offset by the other two.
29
How does the first ionisation energy vary across the period (period 2)
The first ionisation energy increased as we move across a period.
However, in period 2 there are two cases where it decreases .( Beryllium to Boron / Nitrogen to oxygen)
30
Explain how the first ionisation energy generally increases across a period
It tends to increases as we move across a period. The positive charges in the nucleus increases as the number of protons increases. This increases the attraction between the nucleus and the electrons. The atomic radius deceases across a period. Additionally, electron are being added onto the same shell (no increase in atomic radius) and there is similar level of electron shielding (minimal shielding)
-Both the increases nuclear charge and decreased nucleus radius means that the outer electrons are more attracted to the nucleus. This causes the first ionisation energy to increases across a period. (In all these elements, one electron is removed from a shell depending on the group therefore te shielding effect due to the inner electron shell is the same for each element)
31
Exceptions to the increase of first ionisation energy across a period ( period 2)
Boron and Beryllium:
-Beryllium’s outer electron is in the 2s subshell.
-Boron’s outer electron is in the 2p subshell, which has higher energy than 2s.
-Electrons in 2p are less tightly held, so less energy is needed to remove Boron’s electron.
-Therefore, Boron’s first ionisation energy is lower than Beryllium’s.
Oxygen vs. Nitrogen:
-Nitrogen has three unpaired electrons in separate 2p orbitals.
-Oxygen has one 2p orbital with a paired electron.
-Electron-electron repulsion in the paired orbital makes it easier to remove an electron from Oxygen.
-Hence, Oxygen’s first ionisation energy is lower than Nitrogen’s.
32
Explain the pattern of first ionisation energy across a period for period 3
There is a general increase in ionisation energy across the period.
Magnesium and aluminum:
However, there is a decrease in ionisation energy between magnesium and aluminum. This is because the outermost electron of magnesium is on the 3s orbital which has a smaller atomic radius and a greater attraction between the nucleus and outermost electron and less electron shielding. As a result requires more energy to remove the first electron. Whereas, the outermost electron on aluminum is on the 3p orbital which has a larger atomic radius and a weaker attraction between the nucleus and outermost electron and more electron shielding. As a result requires less energy to remove the first electron.
Phosphorus and sulfur:
However, there is a decrease in ionisation energy between phosphorus and sulfur. This is because phosphorus fills it orbital halfway which is more stable compared to sulfur which has a pair of electron in one of the 3p orbitals. This causes more repulsion making it easier to remove the first electron. As a result sulfur requires less ionisation energy to remove the first electron.
33
what do successive ionisation energies tell us how electron are arranged in atoms using oxygen as an example
There is a gradual increase in ionisation energy of oxygen as we remove the first six electrons.
This is because each time we remove an outer electron the remaining electrons in the outer shell are pulled slightly closer to the nucleus. This means that there is a greater attraction between the outer electrons and the nucleus. This causes the ionisation energy to gradually increase.
There is also a massive increase in ionisation energy when the seventh electron is removed. This is explained by looking at the electrons in oxygen. The first six electrons are found in the second electron shell. However, once those electrons are removed the seventh electrons is removed from the first electron shell. Compared to the second electron shell, the first electron shell is closer to the nucleus and electrons in the first shell experience much less shielding.
This means that electrons in the first shell have a greater attraction to the nucleus compared to the electrons in the second shell. This explain why ionisation energy is much greater for the 7th and 8th electron compared to the first and sixth.
34
How to use ionisation energy data to identify an element
You need to find the number of electrons in the outer shell. This can be done through looking at the increase of ionisation energy. (Gradual increase compared to massive increase)
This tells us that the element will have the same number of electrons in the outer shell as from where the gradual increase of ionisation stops
Therefore, the next electron must be removed from an internal shell. Therefore, the ionisation energy is much greater than than the previous ones.
Therefore, we can identify that the number of electrons in the outer shell is the gradual increase from where the ionisation energy stops.
35
Key idea of ionic bonding
Many atoms react in order to to achieve the electron configuration of a noble gas
In ionic bonding, electron are transferred from the metal to the non-metal.
The square brackets around an ion tells us the charge is spread over the whole ion which are attracted to each other through electrostatic forces of attraction
36
What are the key properties of ionic compounds
High b.p/m.p:
They have a very high melting and boiling point. They need a lot of energy to overcome the strong electrostatic forces of attraction between the oppositely charged ions in the lattice.
Solubility:
They are soluble in polar solvents such as water. When we dissolve an ionic compound in water, the water molecule surround the ions which overcomes the strong electrostatic forces of attraction between the oppositely charged ions. Thus many ionic compounds dissolve in polar solvents. ( if charges of ions increase the solubility decreases)
Electrical conductivity:
They do not conduct electricity when solid. The ions are locked in place by the strong electrostatic forces of attraction between oppositely charged ions. Therefore, the ions cannot move from place to place and carry charge. ( if it is melted or dissolved in water it can conduct electricity)
37
When does ionic bonding take place
Transfer of electrons from the metal atom to the non-metal atom
38
When does covalent bonding take place
When a non-metal reacts with a non-metal
39
How does covalent bonding occur
The strong electrostatic forces of attraction between a shared pair of electrons and the nuclei of the bonded atoms.
40
How is a shared pair of electrons represented
A straight line e.g
H —- H
41
What are some exception to non-metal covalently bonding to achieve an electron configuration of a noble gas
-Boron has three electron in its outer shell. Thus often forms a compound with only three covalent bonds.
-Boron Trifluoride has 6 outer shell electrons (does not achieve the same electron configuration as the nearest noble gas)- electron deficient compound
Or
-Phosphorus has 5 electrons in its outer shell. It can form 5 covalent bonds, giving it 10 electrons in its outer shell.
-This is called expansion of the octet — having more than 8 electrons.It happens because period 3 elements have access to the 3d subshell. ( Does not occur in period 1 or 2 as there is no d-subshell)
42
What is a pair of electrons in the outermost shell which is not used to form a covalent bond called
A lone pair of electrons
43
What is a dative covalent bond
When an atom uses a lone pair of electron to form a covalent bond. This is called dative covalent bond.
44
What is an example of dative covalent bond
In ammonia, nitrogen forms three covalent bonds and one lone pair of electrons.
When a hydrogen ion it approaches it cannot contribute any electron to the covalent bond . Therefore, the nitrogen uses its lone pair of electrons to form a covalent bond to the hydrogen ion (dative covalent bond) . This forms the ammonium ion NH4+
45
How is a dative bond shown through displayed formula
A dative bond is shown by an arrow (the arrowhead points away from the element providing the lone pair)
46
Key points of dative bonds
-In order for dative bonds to form, the acceptor must be electron deficient. (Available orbits for electron to occupy)
-The dative covalent bond is exactly the same as normal covalent bond
-All bonds are the same length and have the same average bond enthalpy (tells us the strength of the bond)
47
What do solid lines represent
The solid lines represents that the bond lies on the plane of the screen or page
48
What do solid wedges represent
It shows that the bond is coming out of the plane of the page
49
What does a dotted wedge represent
It shows the bond is going behind the plane of the page
50
Explain EPRT (electron pair repulsion theory)
The shapes of a molecule is determined by the electron pairs surrounding the central atom
We are only referring to the outer shell. This based on the fact that pairs of electrons repel all of the other electron pairs.
The electron pair moves as far apart as possible to minimise repulsion
Dative bonds behave in the same ways as regular covalent bonds
Lone pair reduce the bond angle by an extra 2.5 degrees
51
Key points of EPRT
We treat a multiple bond as a single bonding area
Bond pairs space as far apart as possible
52
What is the bonding angle of linear molecule
180
53
What is the bonding angle of a trigonal planar
120
54
What are the bonding angle of a tetrahedral molecule
109.5
55
How do the bond arrange themselves when they are five pairs of bond angles
Two bonding pair move to opposite sides of the molecule. The other three bonding pairs take up a central position lying on the same plane and spread apart as far as possible
There are two bond angles one pointing up and down the central pane (90) and the angles between the central place is 120
56
what is the shape called when they are five bond angles
Trigonal bipyramidal
Trigonal as the three atoms on the central plane are forming a triangle
Bipyramidal as these form two pyramid shapes with the other two atoms
57
What is the shape called when there are six bonding pairs
Octahedral
58
How do the bond arrange themselves when they are six pairs of bond angles
Two bonding pair move to opposite sides of the molecule. The other four bonding pairs take up a central position lying on the same plane and spread apart as far as possible
There are two bond angles one pointing up and down the central pane (90) and the angles between the central place is 90
59
Effects of lone pairs on the shapes of molecules
Lone pairs repel more strongly than bonding pairs. This extra repulsion decreases other bond angles by 2.5.
60
What is the name of a molecule with a bond angle of 104.5
This is called a non-linear or V-shaped molecule
61
Define electronegativity
Electronegativity is the ability of an atom to attract the pair of electrons in a covalent bond
62
What happens when there is a covalent bond between two atoms of the same elements
Both atoms will have the same electronegativity therefore the pair of electrons in the covalent bond are equally attracted to the two nuclei.
63
What happens when there is a covalent bond between atoms of different elements
-If one atom is more electronegative than the other then this will shift electron density towards the atom leading to the formation of a polar molecule.
-If the molecule is assymetrical than an overall dipole moment will form
64
Where does electronegativity increase in the periodic tale
Electronegativity increase as you go up and right of the periodic table
65
What is electronegativity based on
The Pauling electronegativity scale
66
What does electronegativity depend on
Nuclear charge- The greater the positive charge the greater the attraction between the nucleus and the pair of electrons in the covalent bond.
Atomic radius-The smaller the atomic radius, the closer the bonding electrons will be to the nucleus of an atom
Electron shielding-Shielding of the nucleus by electrons in inner shells. Electron in the inner shell screen electrons in the outer shell from the positive charge of the nucleus. The greater number of inner shells, the lower the electronegativity e
67
What is a pure covalent bond
When the electron pair in the covalent bond lies midway between the two nuclei
68
What do you call when an electron pair is much closer to one atom then the other
This separation of charge is called a dipole moment
69
How can you know a bond is polar
When the delta positive and negative symbols are used to show the charges (delta means the charge is small)
The electron pairs only shift to the more electronegative atom (the delta negative symbol will be on the more electronegative element)
An arrow with a vertical line near the end also shows bond polarity. It points to the more electronegative element.
70
What is a dipole moment
It is the measure of the seperation of charge in a molecule due to difference in electronegativity and the direction of dipoles.
71
What happens to the overall polarity of a molecule if the bonds are in a straight line
Bonds which point in opposite directions in a straight line means that the dipoles cancel each other out. Therefore, the molecule has no overall polarity.
72
Explain the dipole in trichloromethane
The carbon and hydrogen bond have similar electronegativities. However , the dipoles between the carbon to chlorine atoms are polar due to the difference in electronegativity.
Additionally, the dipoles do not cancel as the molecule is asymmetric. As a result there is an overall dipole moment
73
What are the dipoles in the molecule water
Oxygen is mor electronegative compared to hydrogen. Therefore, it shows electron density towards itself. As a result the oxygen atom in water is delta negative and the hydrogen atoms are delta positive. This difference in electronegativity between the atoms makes water polar.
The water molecules has a non-linear shape. This means that although these two bonds point in opposite directions they are not acting in a straight line. The bond polarities cannot cancel out making water polar.
As a result due to the difference in electronegativity and the shape of the bond angle water has a overall permanent dipole
74
What are simple molecular substances
Simple molecular substances consist of relatively small molecules. Each molecule has a fixed number of atoms. These are held together by strong covalent bonds. However, there are weak intermolecular forces of attraction between molecules
75
Explain the properties of simple molecular substances
Boiling/ melting point:
Simple molecular substances are chemically bonded by a covalent bond. They are extremely strong and usually only broken in a chemical reaction. They are not affected during boiling. However, London forces are formed between molecules which are much weaker than covalent bonds and are easily broken (low energy) e.g by high temperature. When we heat a simple molecular substance this causes the molecules to move faster which forces the intermolecular forces to break. This allows the molecules to separate.
Electrical conductivity:
Covalent compound cannot conduct electricity. This is because there are no mobile ions to carry the charge.
Solubility:
Simple molecular lattices that are non-polar dissolve well in non-polar solvents. However, non-polar solvents are insoluble in polar solvents. (Opposite for polar simple molecular lattices)
This is because the hydrogen bonds in polar solvents form hydrogen bonds with each other then formLondon forces with the simple molecular lattice. On the other hand, non-polar solvents interact with simple molecular lattices through London forces resulting in the simple molecular lattice to dissolve
76
How does the bonding of simple molecular substances show the b.p/m.p
Simple molecular substances are chemically bonded by a covalent bond. They are extremely strong and usually only broken in a chemical reaction. They are not affected during boiling.
However, London forces are formed between molecules which are much weaker than covalent bonds and are easily broken (low energy) e.g by high temperature. When we heat a simple molecular substance this causes the molecules to move faster which forces the intermolecular forces to break. This allows the molecules to separate.
Additionally, as simple molecular substances are made up of a fixed number of atoms. The London forces between molecules are weak due to the lack of electrons
77
What are the several types of intermolecular forces
Induced dipole-dipole interactions (London forces / dispersion forces)
Permanent dipole-dipole interactions
Hydrogen bonds
78
Explain London forces
In any atom or molecule, electrons are constantly moving.At any instant, there may be more electrons on one side than the other, creating a temporary uneven distribution of charge. This forms an instantaneous dipole.
The repulsion causes the next atom to have a dipole which is called an induced dipole
All the dipoles will now experience a force of attraction this attraction is called a London force
79
Key points of London forces
London forces are weak and easily broken All intermolecular forces are much weaker than covalent bonds
Once the carbon chain increases there are more electrons resulting in strongly temporary dipole to be formed. As a result the London forces are stronger and so the b/p increases of the molecule increases
Additionally, as the carbon chain increases the surface area of the molecule also increase. Therefore, there is a greater surface area for London forces to form across the molecules. This results in the b/p of the molecule to increase.
London forces are the strongest over short distance. Therefore, London forces may be weaker with molecules with branches
80
How does a permanent dipole-dipole form
When there is a permanent uneven distribution of charge due to a difference in electronegativity between atoms in the bond. There is an an electrostatic forces of attraction between the opposite charges on the molecules which is called a permanent dipole-dipole interaction
Additionally, the dipoles must not cancel out due to the molecules shape
81
Key ideas of permanent dipoles
-Only molecules with a permanent dipole (due to polar bonds and an asymmetrical shape) can experience permanent dipole–dipole forces.
-These molecules have an overall dipole moment — a separation of charge across the molecule.
-In symmetrical molecules, the dipoles cancel out, so there is no overall permanent dipole, and they do not form permanent dipole–dipole interactions.
82
83
Why does boiling point increase going down group 7
-Boiling point increases down Group 7 due to increasing strength of London forces between molecules.
Halogens are simple molecular substances. They exist as diatomic molecules. They are together by weak intermolecular forces — specifically London forces (induced dipole–induced dipole)
-As we go down the group, the number of electrons increases. This causes stronger temporary dipoles to form between molecules. This leads to stronger London forces which require more energy in order to overcome and break. As a result there is a greater b.p down group 7.
84
What is a hydrogen bond
-When a hydrogen atom is covalently bonded to a strongly electronegative element.
-The difference in electronegativity causes an unequal distribution of charge between the molecules.
-As a result the hydrogen becomes partially positive and is attracted to the lone pair of electrons on nearby electronegative atoms
85
How is a hydrogen bond drawn
It must run from the hydrogen directly to the lone pair of electrons (dotted line)
86
Conditions required for hydrogen bonding to take place
A hydrogen atom bonded to a strongly electronegative element (N , O , F)
The electronegative atom must have at least one lone pair of electrons.
87
How do hydrogen bond affect the properties of water
-At low temperature water freezes and turn from liquid to solid.
-Each water molecule forms four stable hydrogen bonds with neighboring water molecules .
-This forms a regular open lattice structure that holds the molecule further apart compared to water.
-This makes ice less dense than water and float
88
What is meant by a simple molecular substance
Simple molecular substances are small and contain a fixed number of atoms that are covalently bonded to each other.
They can form weak intermolecular forces of attraction (London forces) with other molecules of its type
89
Explain why simple molecular lattices have m.p/b.p
Simple molecular lattice:
-Simple molecular substances are chemically bonded by a covalent bond. These covalent bonds are are extremely strong and usually only broken in a chemical reaction.
-However, there are weak intermolecular forces (London forces) between molecules which are easily broken adn do not require much energy in order to overcome.
-Therefore, simple molecular lattice have a low b.p/m.p
90
Explain why simple molecular forces increase in m.p and b.p going down the group
As the number of electron in halogen molecules increases the m.p/b.p increase. This is due to intermolecular forces. The strength of London forces increase with increasing number of electrons.
91
Explain why water has a relatively high m.p/b.p compare to other simple molecular substances
Due to the difference in electronegativity between an oxygen atom and a hydrogen atom electron density shift towards the oxygen atom.As a result the hydrogens have a delta positive charge and the oxygen has a delta negative charge. Therefore, water is polar. Thus are able to form hydrogen bonds between other water molecules which are much stronger compared to London forces and require lots of energy in order to break and overcome .
On the other hand, between simple molecular substances they are joined by weak intermolecular forces of attraction (London forces) which do not require much energy to break. Additionally, as the atoms of simple molecular substances are fixed there is a limited amount of electrons. Therefore,the strength of this intermolecular force of attraction is weak.
92
Explain solubility of non-polar substances in non-polar and polar solvents
The solubility of a simple molecular substance depends on whether the substances is polar or non-polar
Iodine is a good example of a non-polar simple molecular substance. Iodine is a solid at room temperature and pressure. The iodine molecules are held in simple molecular lattice by London forces. Non-polar substances like iodine dissolve very well in non-polar solvents (cyclohexane. The solvent molecule form London forces to the iodine molecules.
In contrast, non-polar substances are insoluble in polar solvents. Water molecules forms hydrogen bonds with each other, therefore if we add iodine to water, the water molecules remain hydrogen bonded to each other rather than forming London forces with iodine molecules. This means iodine is insoluble in water.
93
Explain the solubility of simple molecular substances which are polar
Polar substances dissolve in polar solvents. In solid form, the glucose molecule forms ranges of intermolecular forces with each other. London forces, hydrogen bonds and permanent dipole-dipole interactions will all be acting. If we add glucose to water we will see the same intermolecular forces between the glucose molecules and water molecules
This means that polar substances such as glucose are highly soluble in polar solvent such as water. However, polar substances are usually insoluble in non-polar solvents.
94
Explain the electrical conductivity of simple molecular substances
Simple molecular substances do not contain mobile charged particles.Therefore, simple molecular substances cannot conduct electricity.
95
Explain the electrical conductivity of simple molecular substances
Simple molecular substances do not contain mobile charged particles.Therefore, simple molecular substances cannot conduct electricity.
96
97
What is ionic bonding
Ionic bonding is the electrostatic forces of attraction between positive and negative ions
98
Explain how the solid structures of giant ionic lattices are formed
Giant ionic lattices are formed from the oppositely charged ions which are attracted in all directions by electrostatic forces of attractions.
99
What are the conditions required for a molecule to have an overall dipole moment
Covalently bonded atoms with different electronegativities
A polar molecules requires polar bonds with dipoles that do not cancel due to their direction
100
What are the solid structure of simple molecular lattices
Covalently bonded molecules with a small number of atoms.
Theses molecules are attracted by relatively weak intermolecular forces of attraction (London forces) between molecules
101
What is average bond enthalpy referred as
As a measurement of the covalent bond strength
102
What is the name of the shape and bond angle for:
1) 2 bonding pairs
2) 3 bonding pairs
3) 2 bonding pairs and a lone pair
4) 4 bonding pairs
5) 3 bonding pairs and a lone pair
6) 2 bonding pairs and a 2 lone pair
7) 5 bonding pairs
8) 4 bonding pairs and a lone pair
9) 3 bonding pairs and 2 lone pairs
10) 6 bonding pairs
11) 5 bonding pairs and a lone pair
12) 4 bonding pairs and 2 lone pairs
1) Linear (180)
2) trigonal planar (120)
3) Bent (117.5)
4) Tetrahedral (109.5)
5) Trigonal pyramidal (107)
6) Bent (104.5)
7) Trigonal bipyramidal (120 along the plane 90 above and below)
8) Trigonal bipyramidal (119 along the plane 89 above and below)
9) Trigonal bipyramidal (120 along the plane 89 above and below)
10) Octahedral (90)
11) square pyramid (89)
12) Square planar (90)
