Electrons, bonding and structure Flashcards
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.
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, it only has one electron in the 4s subshell, it can have a half full 3d subshell
In the case of copper, by only having one electron in the 4s subshell , it can have a full 3d subshell
14
Q
How is the periodic table divided
A
The periodic table is divided into 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
Complications of the d-block in the periodic table
A
The first row of the d-block represents the electrons in the d-sub-shell of the third electron shell.
The 4s subshell fils before the 3d subshell
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 nit 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 molecule 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 measuring the ionisation energy each time
25
Differences between first and second ionisation energy
Second ionisation energy removes one molecule of electrons from one mole of 1+ ions not one mole of atoms
We are also making one mole of 2+ ions
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
As we move across a period, the nuclear charge increases as the number of protons increases.
This increases the attraction between the nucleus and the electrons. Therefore, the atomic radius decreases across a period.
Both the increases nuclear charge and the decreased atomic radius means that the outer electrons are more attracted to the nucleus.
As a result the first ionisation energy increases across the period
In every element we remove an electron from the same electron shell. This means that the shielding effect due to the inner electron shell is similar for each element.
31
Exceptions to the increase of first ionisation energy across a period ( period 2)
Boron and oxygen do not fit the pattern of increasing first ionisation energy. To explain this we look at the subshells involved.
In the case of beryllium and lithium we remove an electron from the 2s subshell.
In Boron, the outer electrons is now in the 2p subshell which has a higher energy than the 2s subshell. Thus less energy is needed to remove the outer electrons of Boron compared to the outer electron of Beryllium. Therefore Boron has a lower first ionisation energy than Beryllium.
First ionisation energy increases across carbon and nitrogen but then falls at oxygen. To explain we then look at the 2p subshell.
Nitrogen has each electron in a separate 2p orbital whereas in oxygen one orbital has a pair of electrons.
These electrons repel which mean that it takes less energy to remove one of these electrons than if they were in separate orbitals. Therefore the first ionisation energy of oxygen is less than nitrogen.
32
Explain the pattern of first ionisation energy across a period for period 3
At first there is a general increase in first ionisation energy but a decrease from magnesium to aluminum / phosphorus to sulfur. The pattern of change is the similar to period 2. However, the third energy shell is affected (in period 3) not the second during period 2
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
They have a very high melting and boiling point. They need a lot of energy to overcome the strong electrostatic forces of attraction.
They are soluble in polar solvents such as water. When we dissolve an ionic compound in water, which is a polar molecule the water molecule surround the ion which overcomes the electrostatic forces of attraction between ions thus many ionic compounds dissolve in polar solvents. ( if charges of ions increase the solubility decreases)
They do not conduct electricity when solid. The ions are locked in place by the electrostatic forces of attraction. 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 pair of electrons is attracted to the two nuclei of the atoms forming a bond. Both atoms now have a full shell of electron in the outer shell
40
How is a shared pair of electrons represented
A straight line e.g
H —- H
41
What are some exception to non-metal covalent bonded to achieve an electron configuration of a noble gas
Boron has three electron in its outer shell. Thus often forms compound with only three covalent bonds.
E.g Boron Trifluoride has 6 outer shell electrons (does not achieve the same electron configuration as the nearest noble gas)- electron deficient compound
Or
E.g :Phosphorus has 5 electron in the outermost shell. Therefore forms 5 covalent bond to achieve an outer shell of 10 electrons. Therefore, not achieving the electron configuration of the nearest noble gas (this is because phosphorus is in period 3 and has the 3d subshell)- expansion of the octet (cannot take place with element in period 1 or 2 as these do not have a 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 there is a hydrogen ion 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)
The dative bond has the same bond strength as the other bonds
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 projecting back 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
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
The two different elements will have different levels of electronegativity. Therefore, the atom with a greater electronegativity will have the electron pair in the covalent bond closer to that nucleus.
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
The size of the positive charge on the nucleus. The greater the positive charge the greater the attraction between the nucleus and the pair of electrons in the covalent bond.
The smaller the atomic radius, the closer the bonding electrons will be to the nucleus of an atom
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
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
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
When the only bond is polar thus has an overall polarity.
71
What happens to the dipoles with molecules with more than one bond
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
What are the dipoles in the molecule trichloromethane
The carbon and hydrogen bond have similar electronegativities. Therefore, the dipoles on the carbon to chlorine bonds cannot cancel , this is a polar molecule.
The side with hydrogen is positive and the side with chlorine is negative.
73
What are the dipoles in the molecule water
water contains two hydrogen atoms bonded to one oxygen atom. These bonds are 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
74
What are simple molecular substances
Molecules which consist of of relatively small molecules. Each molecule has a fixed number of atoms.
75
Properties of simple molecular substances
They have low b.p and m.p
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, intermolecular forces 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.
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 an atom or molecule electrons are randomly moving around. If there are more electrons on one side there is a dipole (negative charge on the side with more electrons and positive with the side with less). This is an instantaneous dipole as it happens in a particular instant.
The repulsion causes the next atom to have a dipole which is called an induced dipole (“induced” caused by something else). Thus the dipole did not happen randomly.
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
London forces are caused by random electron movement (every single atom or molecular experiences London forces) even if they experience other intermolecular forces
The strength of London forces depends on the number of electrons
80
How does a permanent dipole-dipole form
When two molecules with a permanent dipole is attracted to each other the permanent dipoles can lead to an attraction called a permanent dipole-dipole interaction. (Intermolecular force)
81
Key ideas of permanent dipoles
Only molecules with a permanent dipole can experience this type of intermolecular force.
Molecule with a permanent overall dipole (dipole moment) can form a permanent dipole-dipole interaction. On the other hand, symmetrical molecules would cancel out their polar bonds and have no overall permanent dipole.
82
Explain why does trichloromethane which has permanent dipole-dipole interactions not have a higher melting point then tetrachloromethane
Permanent dipole-dipole interaction are not only the forces acting. All molecules and atoms experience London forces.
The size of London forces depend on electrons present. 58 electron in trichloromethane whilst tetrachloromethane has 74. This means that London forces are stronger in tetrachloromethane. Therefore, has a higher b.p
83
Why does boiling point increase going down group 7
All hydrogen halides have a permanent dipole so permanent dipole-dipole interactions are acting in all three cases.
The strength of the permanent dipole decreases going down the table. That’s due to the decreasing electronegativity of the halogens as we move down group 7. Therefore the m.p/b.p decreases.
However, the boiling point increases as you go down group 7 as the molecules of he hydrogen halides increase in number of electrons. Therefore , the London forces increase as we make our way down the table. This causes the boiling point of hydrogen halides to increase
84
What is a hydrogen bond
When a positive hydrogen atom is attracted to the lone pair of electrons on another molecule.
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.
The electronegative atom must have at least one lone pair of electrons.
87
How do hydrogen bond affect the properties of water
when water is boil it takes a greater deal of energy to break the hydrogen bonds. Therefore, water has a relatively high boiling point of 100c. (Along with a relatively high m.p)
Ice is less dense than liquid and floats on the surface of the water. This is due to hydrogen bonding. In liquid water, the water molecules move randomly. Sometimes they are close together or far apart. Hydrogen bonds are constantly being formed and broke. As we cool the water down, the molecules move more slowly As we react the freezing point (0c) the water molecules arrange themselves into an ordered structure.This is ice which is stabilized by the network of hydrogen bonds. In ice, the water molecules are further apart than in liquid water. This makes ice less dense and float on water. Ice acts as an insulating layer for organism below.
88
What is meant by a simple molecular substance
In simple molecular substances the atoms are covalently bonded to each other. Secondly, simple molecular substances have small molecules with a fixed number of atoms.
89
Explain why simple molecular forces have a low m.p/b.p
They have low melting and boiling points due to their intermolecular forces. If we cool any simple molecular substance to below its melting point then it will form a solid. This is called a simple molecular lattice.
In halogens molecules consist of two atoms of the same type joined by a single covalent bond. In between the molecule is intermolecular forces.
The molecules consist of the same atom therefore have the same electronegativity thus will be non-polar. This means that only induced dipole-dipole interaction (London forces) will act between the molecules.London forces are weak and take little energy too break so because simple molecular substances have weak intermolecular forces they all have low m.p/b.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
Water also has a relatively high melting and boiling point compared to other simple molecular substances. In ice the water molecules are held in place by intermolecular forces. These intermolecular forces include hydrogen bonds. When ice is melted these hydrogen bonds are broken, because hydrogen bonds are relatively strong compared to other intermolecular forces it takes quite a lot of energy to break hydrogen bonds.
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 how simple molecular substances which are polar react
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
An electrical current is a flow of charged particles. In metal, the charged particles are delocalised electrons. whereas, in a molten ionic compound, the charged particles are mobile ions.
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
An electrical current is a flow of charged particles. In metal, the charged particles are delocalised electrons. whereas, in a molten ionic compound, the charged particles are mobile ions.
Simple molecular substances do not contain mobile charged particles.Therefore, simple molecular substances cannot conduct electricity.
96
