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A Level Chemistry > Atomic Structure - Topic 1 > Flashcards

Atomic Structure - Topic 1 Flashcards

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1
Q

How many orbitals and electrons in the s orbital

A

1 orbital
2 electrons

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2
Q

How many orbitals and electrons in the p orbital

A

3 orbitals
6 electrons

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3
Q

How many orbitals and electrons in the d orbital

A

5 orbitals
10 electrons

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4
Q

How many orbitals and electrons in the f orbital

A

7 orbitals
14 electrons

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5
Q

What are the two isotopes of chlorine

A

35Cl 37Cl

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6
Q

What blocks are group 1 and 2 and why

A

S blocks
As there outer electrons are in energy levels called s subshells

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7
Q

What blocks are Group 3 and 0 found in and why

A

P blocks
As there outer electrons are in energy levels called p subshells
(Except helium)

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8
Q

Properties of isotopes

A
  • identical chemical properties since they have the same electron configuration
  • different physical properties
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9
Q

What does a mass spectrometer do

A

Helps as determine the abundance and molecular mass of isotopes

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10
Q

The stages of mass spectrometry

A

1) vapourastion
2) ionisation
3) acceleration
4) deflection

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11
Q

How are the atoms ionized when in a mass spectrometer

A

Atom are turned into its gaseous state then are ionized in the spectrometer by high energy electrons bombarding the sample, knocking off their electrons forming a positive ion and a radical ( a single electron)

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12
Q

What happens in the acceleration and deflection stage in a mass spectrometer

A

1) ions are accelerated by an electric field, so they have the same KE
2) ions stop accelerating then drift to a magnetic field where they are deflected

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13
Q

How does the mass spectrometer discover the masses of isotopes

A

By recording the time it takes for the ions to deflect off of the magnetic field as the lighter ions would drift at a higher velocity so would deflect first compared to the heavier ions

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14
Q

How does the mass spectrometer determine the abundance of an isotope

A

When the ions reach the magnetic field it gains an electron generating a current and the size of the current determines the abundance as it give a measure of the number of ions being deflected

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15
Q

Ratio of 35Cl to 37Cl

A

3:1

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16
Q

Ratio of 35Cl-35Cl to 37Cl-35Cl or 35Cl-37 to 37Cl-37Cl

A

9:6:1

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17
Q

Abundance 35Cl

A

75%

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18
Q

Abundance of 37Cl

A

25%

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19
Q

Lithium isotopes

A

6Li 7Li 8Li

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20
Q

Mass of electron proton and neutron

A

Electron 1/1840
Proton 1
Neutron 1

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21
Q

Relative intensity for 79Br 81Br 158Br 160Br and 162Br

A

79Br - 10
81Br - 10
158Br - 50
160Br - 100
162Br - 50

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22
Q

Ratio for 79Br to 81Br

A

1:1

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23
Q

Ratio for 158Br to 160Br to 162Br

A

1:2:1

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24
Q

Isotopes of bromine

A

79Br 81Br

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25
What is the principal quantum number,n
The number given to each shell
26
How to find the maximum number of electrons found in each shell
2n^2
27
What do atomic orbitals show
A 95% probability of where an electron would exist
28
What are the atomic orbitals
S P D F
29
First ionisation energy
The energy needed to remove 1 mole of electrons from one mole of an atom in its gaseous state to form 1 mole of 1+ ions in their gaseous state
30
Second ionisation energy
The energy needed to remove 1 mole of electrons from one mole of 1+ ions in their gaseous state to form 1 mole of 2+ ions in their gaseous state
31
Equation for first ionisation energy
X(g) ———> X^+(g) + e-
32
Factors that affect ionisation energy
- atomic radius - shielding - charge on the nucleus
33
How does atomic radius affect ionisation energy
As the atomic radius increases the force of attraction between the positive nucleus and outer electron decreases so ionisation energy decreases
34
How does charge of the nucleus affect ionisation energy (IE)
The greater the number of protons the greater the force of attraction between the outer electrons and the nucleus so IE increases
35
How does shielding of the nucleus affect ionisation energy (IE)
Electrons in the outer shell are repelled by electrons in the inner shell so the more shielding the smaller the attraction between the outer electron and nucleus so the IE decreases
36
Why is their a gradual increase in ionisation energy on an ionisation energy graph
As when we remove a valence electron the remaining electrons in the outer shell is pulled closer to the nucleus so there is a greater attraction between the outer electrons and the nucleus causing the IE to gradually increase and all electrons are being removed from the same quantum shell
37
What does the gradual increase in IE repersent on an ionisation graph
The number of electrons in the outer shell
38
What does the massive increase in IE show on a ionisation graph
That the electron that we are removing is in a shell closer to the nucleus than the previous electrons and it sits alone in the outer most shell so it requires more energy to remove
39
Why does the first ionisation decrease as you go down a group
- Although the number of protons increase it is outweighed by the shielding and atomic radius - As the shielding increase there is more electron repulsion between the outer electrons and the inner electrons causing the ionisation energy to decrease , the increase in shielding also causes the atomic radius to increase as the electrons are becoming further away from the nucleus so the attraction decreases
40
Why does the first ionisation increase as you go across a period
- the positive charge in the nucleus increases as the number of protons increases, increasing the attraction between the nucleus and electrons, causing the atomic radius to decrease - the increase in nuclear charge and decrease in atomic radius increases the attraction between the outer electrons and the nucleus causing the ionisation energy to increase
41
why can the first ionisation energy DECREASE as you go across a period of number of electrons are the same
- when removing an electron from a higher energy level the energy level is further from the nucleus so it takes less energy to remove the outer electron so the IE decreases - if they are in the same orbital the sub shell with the paired electrons would require less IE to be removed as they repel each other so take less energy to remove than if they were in separate orbitals
42
How are emissions spectrums produced
When energy is emitted by excited electrons dropping to lower energy levels
43
What are the lines in the visible spectrum caused by
Electrons dropping to n=2 energy level
44
What are the lines in the UV part of the spectrum caused by
Electrons dropping to n=1 energy level
45
What are the lines in the infrared part of the spectrum caused by
Electrons dropping to n=3 energy level
46
What happens when you pass a lot of energy through an hydrogen atom
The electron only absorbs the exact energy difference from the first energy level to the second and moves to the higher energy level (and when it drops it releases the exact same amount of energy due to the conservation of energy)
47
What is hunds rule
the electrons will occupy the orbitals singly before pairing takes place
48
What are the periodic trends
Atomic radius Ionisation energy Electron affinity Electronegativity Bonding structure (mp and bp)
49
What is the atomic radius
The distance from the nucleus to the electrons in the outer energy level
50
Why does atomic radius decrease across a period
- as we move across a period the proton number increased by 1 so the positive charge increases this increase in positive charge means that there is a greater attraction between the nucleus and electrons drawing the electrons closer to the nucleus causing the atomic radius to decrease - shielding due to the inner electrons is the same across the period
51
Why does atomic radius increase moving down a group
The number of electron shells increase as you go down a group this is because each element has one more full inner electron shell increasing the amount of shielding between the nucleus and outer electrons so the outer electron shell is further away from the nucleus causing there to be less attraction between the outer electrons and the nucleus and increasing atomic radius
52
Bonding and structure - metallic bonding
- the negatively charge delocalised electrons are strongly attracted to the positive metal ions by strong metallic bonding - this overall structure is called a giant metallic bond
53
Key feature of giant metallic lattice
- cations (metal ions) are fixed in place and cannot move - delocalised electrons are free to move so can conduct heat and electricity when solid and liquid - have high melting point and boiling point due to strong metallic bond - don’t dissolve
54
Bonding and structure - property of giant covalent structures
- have high melting and boiling points as billions of atoms are joined together by strong covalent bonds to form a giant covalent lattice and it would require lots of energy to overcome these bonds
55
Key features of giant covalent bonds
- high melting and boiling point - does not conduct electricity as it has not delocalised electrons or ions since every electron is involved in covalent bonding - insoluble as solvents cannot disrupt the large number of strong covalent bonds
56
Why can graphite conduct electricity even though it’s a giant covalent bond
As the carbon atoms form 3 covalent bonds instead of 4 (like in diamond) so it has one delocalised electron
57
Bonding and structure - simple molecular bonding
They have a simple molecule lattice with weak intermolecular forces between molecules which do not require lots of energy to break so have low bp and mp
58
Examples of giant covalent structures
Boron Carbon Silicone
59
What is electron affinity
How easy it is to gain an electron
60
What happens to the electron affinity across a period
It increases
61
Why does the ionisation energy increase each successive ionisation energies
As when an electron is removed from an ion or an atom the remaining electrons experience a greater form of attraction to the nucleus so the ionisation energy increases
62
know how to draw ionisation energy graph
63
why can some atoms have more the 8 valence electrons whilst others can't
- as for elements period three and beyond they can expand their octet so can accommodate for more than 8 valence electrons as they have empty d orbitals so can use their d orbitals to accommodate more than 8 electrons - however for period 2 elements there are no available d orbitals for expansion, and their small size leads to significant electron-electron repulsion. Therefore, they are cannot accommodate for more than 8 electrons

A Level Chemistry (23 decks)

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