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(A-level) Chemistry > Module 2.2 > Flashcards

Module 2.2 Flashcards

(96 cards)

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1
Q

What happens when atoms lose /gain electrons

A

They become ionised (positively if they gain, negatively if they lose)

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2
Q

How are ionising patterns used to tell the number of shells an atom has

A

Energy is applied to an atom, exciting an electron, so that it leaves a shell

Electrons in the same shell are excited at similar energy levels

Where there is a large increase in energy there is a new shell

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3
Q

Why is more energy needed each time to remove an electron from an atom

A

When you remove an electron from the outer shell there is a greater force (applied by the nucleus) on the remaining electrons, so more energy is needed

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4
Q

What is an electron shell

A

The energy level which an electron orbits the nucleus

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5
Q

What are orbitals

A

The direction / orbit an electron follows at its energy level

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6
Q

What are S orbitals

A

The first orbital in the shell
They are circular and hold 2 electrons

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7
Q

How many electrons can each orbital hold

A

2

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8
Q

What are P orbitals

A

There are 3 P orbitals which all move in a figure 8

Px - along the x axis
Py - along the y axis
Pz - along the z axis

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9
Q

Why is there 1 S orbital and 3 P orbitals in the 2nd shell

A

As they hold a total of 8 electrons

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10
Q

Write the orbitals for sodium (NA11)

A

1S2 /\ 2S2 2Px2 2Py2 2Pz2 /\ 3S1

1st number = shell
2nd number = number of electrons

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11
Q

In what order do the p orbitals fill

A

Each orbital must be half full before one can hold 2 electrons

Px then Py then Pz

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12
Q

Why does hydrogen share electrons with all elements (orbitals)

A

Hydrogen has one electron so a 1S1 orbital

All orbitals / elements want to have full shells and become neutral, so it shares this electron with another

for example 1 hydrogen would bond to 1 sodium

The electron can be shared to all orbitals

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13
Q

What is shorthand notation (Aufbau principle)

A

A notation used to shorten writing out all the orbitals of an atom
It uses a noble gas as the next orbital will always be an S orbital

E.G Rubidium

[Kr] 5S1

where Kr is Krypton

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14
Q

How many electrons (in total) do D orbitals contain

A

10 - so 5 individual orbitals

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15
Q

How many electrons (in total) do F orbitals contain

A

14 - so 7 individual orbitals

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16
Q

When do D orbitals first appear

A

shell 3

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17
Q

When do F orbitals first appear

A

shell 4

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18
Q

How many electrons are in shell:
1
2
3
4
5

A

1 = 2 electrons
2 = 8 electrons
3 = 18 electrons
4 = 32 electrons
5 = 32 electrons

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19
Q

In what order to the orbitals fill

A

1s - 2s - 2p - 3s -3p -4s -3d - 4p - 5s - 4d - 5p - 6s - 4f - 5d - 6p - 7s

Fills in a snake pattern, the shell bellows s orbital will fill before the previous shells d orbitals

1s
2s 2p
3s 3p 3d
4s 4p 4d 5f
5s 5p 5d 5f

etc………………..

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20
Q

What is the difference between an orbital and a subshell

A

A subshell is the pathway in a shell which an electron moves. (the number of electrons in a subshell depends on the subshell)

An orbital is a mathematical function that shows the wave like nature of an electron. (orbitals can only have 2 electrons)

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21
Q

How does the size of atoms change moving from metals to non metals

A

Moving from the left of the periodic table to the right the atomic radius decreases

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22
Q

Why does the atomic radius decreases moving towards the non metals

A

The number of protons in the nucleus increases, so the nucleus has a greater NUCLEAR ATTRACTION to each electron

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23
Q

Why does the ionization energy needed increase moving left to right in the periodic table

A

The electrons require more energy to be removed from the orbitals

except in group 3 and 6

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24
Q

Considering Period 2, why is the ionization energy of group 3 (boron) an anomaly

A

Boron requires less energy to remove an electron from its outer subshell compared to Beryllium

Be = 1s2

B= 1s2 1p1

It requires less energy to remove an electron from the unstable p orbital, than the complete s subshell.

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25
Considering Period 2, why is the ionisation energy of group 6 (oxygen) an anomaly
Oxygen requires less energy to remove 1 electron from its outer subshell than Nitrogen N = 1s2 1px1 1py1 1pz1 O = 1s2 1px2 1py1 1pz1 IT requires less energy to remove an electron from the unstable 1px2 orbital in oxygen that to remove 1 electron from the half filled stable p subshell in nitrogen
26
When are subshells stable
When they are fully filled When they are half filled (1 electron in each orbital)
27
Why does it require more energy to remove electrons from stable subshells than unstable subshells (with 1 additional electron)
Orbitals in stable subshells want to remain stable Orbitals in unstable subshells of +1 electrons want to lose that electron to become stable. As they want to lose the electron it requires less energy
28
What is charge density
The ratio of an ions charge to its volume
29
Which has a greater charge density Sodium or aluminum ions
Na + is in group 1 Al 3+ is in group 3 Aluminum has a smaller atomic radius and a greater charge, so its charge density is greater
30
How does the force of attraction between ions change with charge density
The greater the charge density the greater an ions electrostatic force of attraction
31
How do metals and non metals bond
They form ionic bonds Metals donate electrons to have a full outer shell Non-metals receive electrons to have a full outer shell This makes the atoms oppositely charged so they attract each other
32
What are the general properties of ionic compounds
They have a giant structure - so they also have a high melting and boiling point They can conduct electricity when dissolved or molten (as they contain charged particles - the ions) They are soluble in polar solvents
33
What is the structure of a giant ionic lattice
Each ion is surrounded by an oppositly charged ion The ions attract each other forming a giant ionic lattice
34
What determines the melting and boiling point in ionic compunds
The stronger the electrostatic force of attraction between ions the higher the melting / boiling point will be, as more energy is needed to break the ionic lattice The strength of the attraction is determined by the ions charge density
35
Why cant ionic compounds conduct electricity when solid, but can when molten / dissolved
When solid, the ions are held in a fixed position and cannot move When molten or dissolved the solid lattice breaks down, and ions can move As the ions can move they can conduct electricity
36
What substances can ionic lattices dissolve in
Polar substances such as water
37
how do ionic lattices dissolve in water
Water surrounds the ionic lattice, and it begins to break down The partially positive hydrogen bonds to the negative ion The partially negative oxygen binds to the positive ion
38
What is a covalent bond
A bond formed when two non-metals share electrons
39
How are electrons shared in covalent bonds
The negative electrons are attracted to the nuclei of both atoms The attraction of these electrons overcomes the repulsion force between the two positive nuclei This forms a covalent bond where electrons are shared
40
How are covalent bonds drawn
Using dot and cross diagrams Only the outer shell is drawn The shared electrons are drawn in the overlapping space between the two electrons
41
What is a single bond
The covalent bond where only 1 pair of electrons are shared
42
What are lone pairs
A concentrated space / region of negative charge around an atom 2 electrons that are not used in bonding form a lone pair
43
What orbital is most commonly a lone pair in covalent bonding
The s orbital This is because it is typically full, so electrons do not need to be shared with it in order to make the orbital stable
44
How are double bonds (covalent) formed
When the atoms share 2 pairs of electrons
45
How are triple bonds (covalent) formed
When atoms share 3 pairs of electrons
46
What type of reaction is bond breaking
exothermic
47
Which type of covalent bonds have the highest average bond enthlapy
Triple bonds will require a greater change in enthlapy to break than double or single bond
48
What is Dative covalent bonding
A bond formed where one atom provides both of the shared electrons in a covalent bond
49
How is a dative bond written
A (arrow) B The direction of the arrow shows the direction the pair of electrons are donated
50
What type of atoms can form dative bonds
Dative bonds form between an atom with at least one lone pair and an atom (typically a positive ion) that has an electronic configuration of an empty orbital
51
How many covalent bonds can an atom form
- All unpaired electrons can pair up between atoms - The maximum number of atoms that can pair up is equivelent to the number of electrons in the outer shell
52
What is an octet
a full outer shell
53
What is the octet rule
The octet rule states that: All atoms prefer to have 8 electrons in their outer (valence) shell. This give them a stable electronic configuration
54
What are valence electrons
Electrons in an outer shell
55
Why does expansion of the octet occur
Moving down the periodic table, more of the outer shell electrons are able to take part in bonding elements from group 5 to group 7 (g15 to g17) can expand their octet
56
Why do some atoms not complete their octet in covalent bonding
Atoms such as Boron do not have enough unpaired electrons They have 3 electrons in an outer shell, so can have a total of 6 electrons once paired
57
How many covalent bonds can a group 5 element form
3 or 5 covalent bonds
58
How many covalent bonds can a group 6 element form
2, 4 or 6 covalent bonds
59
How many covalent bonds can a group 7 element form
1, 3, 5 or 7 covalent bonds
60
What are the two types of covalent structure
Simple molecular lattice Giant covalent lattice
61
What are simple molecular lattices and their properties
MAde from small simple molecules such as H2O The atoms in each molecule are held together by covalent bonds The different molecules are held together by weak intermolecular forces They have low melting and boiling points (weak intermolecular forces) They dont conduct electricity They are soluble in non-polar substances
62
Why are simple molecular structures / lattices soluble in non-polar solvents
Weak london forces can form between covalent molecules and non-polar solvents This causes the molecular lattice to break down, dissolving the substance
63
What are giant covalent lattices and their properties
Lage repeating structures such as diamond - where there are covalent bonds between all atoms They have high melting and boiling points (due to strong covalent bonds) They cannot conduct electricity as there are no free charged particles (except in graphite) They are completely insoluble as the covalent bonds in the lattice are too strong to be broken by polar or non-polar substances
64
What types of replusion are there between atoms that are covalently bonded
Bonding pair - bonding pair repulsions If lone pairs are present: Lone pair - bonding pair repulsion Lone pair - lone pair repulsion Bonding pair repulsions are the weakest
65
Why are there bonding pair repulsions in covalent bonds ***********************
When sharing electrons, the atoms become charged and repel each other The share of electrons also repel each other
66
How do lone pairs repel bonding atoms in covalent bonds
They repel atoms downwards giving them a 3D structure
67
How many bonds / what is the bond angle in a *linear(1)* shaped covalent bond
There is 1 covalent bond (so there is no bond angle)
68
How many bonds / what is the bond angle in a *linear(2)* shaped covalent bond
There are 2 covalent bonds There is a 180 bonding angle between bonded atoms
69
How many bonds / what is the bond angle in a *trigonal planar* shaped covalent bond
There are 3 covalent bonds There is a 120 bonding angle between bonded atoms
70
How many bonds / what is the bond angle in a *tetrahedral* shaped covalent bond
There are 4 covalent bonds formed There is a 109.5 bonding angle between bondinded atoms
71
How many bonds / what is the bond angle in a *trigonal bipyramid* shaped covalent bond
There are 5 covalent bonds There is a 90 bonding angle between the bonded atoms in the plane There is a 120 bonding angle between bonded angles coming out of the plane (towards and away)
72
How many bonds / what is the bond angle in a *octahedral* shaped covalent bond
There are 6 covalent bonds There is a 90 bond angle between the bond in the plane and the bond coming out of the plane
73
How many bonds / what is the bond angle in a *pyramidal* shaped covalent bond (And lone pairs)
3 covalent bonds 1 lone pair 107° between each bonding pair
74
How many bonds / what is the bond angle in a *non-linear* shaped covalent bond (And how many lone pairs)
2 covalent bonds 2 lone pairs 104.5° bonding angle
75
When drawring covalent bonds what does the coloured in 'wedge'mean
The bond is coming out of the plane (towards you)
76
When drawring covalent bonds what does the dotted 'wedge'mean
The bond is going out of the plane (away from you)
77
What is electronegativity
The ability of a non-metal in a covalent bond to draw electrons to itself
78
What atom has the greatest electronegativity
Fluorine
79
What is the general trend of electronegativity
Moving up the periodic table and to the left, elements become more electronegative
80
What is the general trend of electronegativity
Moving up the periodic table and to the left, elements become more electronegative
81
What is a polar covalent bond
When the bonding atoms are different, one will have a greater electronegativity THE BONDING PAIR of electrons are attracted more to the more electronegative element This creates a small charge difference 1 atom has a partially positive charge 1 atom has a partially negative charge
82
How do non polar molecules contain polar bonds
If an atom is symmetrical the dipoles can cancel out (as they act in different directions) This makes the atom non polar
83
What is a permanent dipole
A small charge difference across a bond (due to the electronegativity of atoms)
84
What are van der Waal forces
Intermolecular forces including: London forces Permanent polarity (due to electronegativity) Hydrogen bonds
85
What is permanent polarity
A type of intermolecular force where one atom in a covalent bond is partially negative, and the other is partially negative (Due to electronegativity)
86
What are london forces
An intermolecular force caused by the constant and random movement of electrons in an atoms shell
87
How do london forces form
Electrons are constantly moving around an atom / molecule At any time, one side might contain more electrons than the other (making it partially negative) This creates an **instantaneous dipole** across the molecule The **instantaneous dipole** induces a dipole in the neighbouring molecule (which then induces more dipoles) These induced dipoles attract each other creating weak intermolecular forces known as **London Forces**
88
Are london forces stronger or weaker when there are more electrons
The more electrons there are, the larger the induced dipole, the larger the attractive force between molecules
89
What are hydrogen bonds
The strongest intermolecular forces. They occur in molecules containing O-H N-H F-H Hydrogen bonds occur between molecules which contain these bonds
90
Why are hydrogen bonds the strongest intermolecular force
Oxygen, nitrogen and fluorine are very electronegative elements This means elements with these bonds have strong polar dipoles
91
How do hydrogen bonds affect the properties of water
It causes water to have a higher boiling point than expected It causes water to be less dense in its solid state
92
Why is ice less dense than water
When ice forms water molecules arrange in an orderly pattern, allowing hydrogen bonds to form This creates an open lattice (with hydrogen bonds holding the molecules apart) Due to the empty space created, ice is less dense than watet
93
What are ionic bonds
Bonds formed by Electrostatic attraction between positive and negative ions
94
Describe the structure of water
Water is made of a central Oxygen covalently bonded to two hydrogen atoms The oxygen is partially negative and the hydrogen is partially positive The oxygen has two lone pairs It forms a shape with a 104.5° bonding angle as there is lone pair - lone pair repulsion, bonding pair - bonding pair repulsion and lone pair - bonding pair repulsion, where the lone pairs repel more than the bomding pairs
95
What bond is in N2
A covalent triple bond
96
How does electronegativity affect bond enthalpy
If the atoms have a similar electronegativity the enthalpy is higher

(A-level) Chemistry (9 decks)

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