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Chemistry > TOPIC 1: ATOMIC STRUCTURE & THE PERIODIC TABLE > Flashcards

TOPIC 1: ATOMIC STRUCTURE & THE PERIODIC TABLE Flashcards

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1
Q

Mass Number

A

the total number of protons and neutrons in the nucleus

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2
Q

Atomic (proton) number

A

number of protons in the nucleus

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3
Q

Isotope

A

Atoms of the same element with the same number of protons but different number of neutrons.

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4
Q

Describe the Chemical Properties of an Isotope compared to the Element

A

The chemical properties of an isotope are the same as the element as they have the same electron configuration

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5
Q

Describe the Physical Properties of an Isotope compared to the Element

A

The physical properties are slightly different compared to the element as they have different mass numbers

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6
Q

Relative Isotopic mass

A

The mass of an atom of an isotope compared with 1/twelfth the mass of a carbon-12 atom

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7
Q

Relative Atomic Mass

A

The average mass of an atom compared to 1/twelfth the mass of a Carbon-12 atom

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8
Q

The Relative Molecular Mass (when is it used)

A

Used for simple molecules

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9
Q

Define Relative Molecular Mass

A

The average mass of a molecule compared to 1/twelfth the mass of a carbon-12 atom

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10
Q

Define Relative Formula Mass

A

The average mass of a formula unit compared to 1/twelfth the mass of a carbon-12 atom.

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11
Q

What does the Molecular ion peak tell us?

A

The relative mass of a compound

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12
Q

Maximum Number of Electrons in an S subshell

A

2 electrons

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13
Q

Maximum Number of Electrons in a P subshell

A

6 electrons

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14
Q

Maximum Number of Electrons in a D subshell

A

10 electrons

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15
Q

Maximum Number of Electrons in an F subshell

A

14 electrons

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16
Q

First Ionisation Energy

A

The ionisation energy is the energy needed to remove 1 mole of electrons from one mole of gaseous atoms to form one mole of positive ions.

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17
Q

Three Factors that effect Ionisation Energy

A

nuclear charge, Electron shell, Shielding

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18
Q

Periodicity

A

The repeating trend of chemical and physical properties of the elements in the periodic table

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19
Q

what affects atomic radius?

A
  1. Number of protons

2. Number of Shells

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20
Q

Isoelectronic

A

ions of different elements having the same number of electrons

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21
Q

Name giant covalent structures

A
  1. Silicon dioxide
  2. Diamond
  3. Graphite
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22
Q

When is the relative formula mass used

A

Used for ionic or giant covalent compounds

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23
Q

Describe the process of mass spectrometer

A
  1. Ionisation

The atom is ionised by using an electron gun that knocks off electrons giving a positive ion

  1. Acceleration

The ions are accelerated by an electric field so that they have all the same kinetic energy

  1. Deflection

The ions are then deflected by a magnetic field with the amount of deflection depending on their masses. The lighter they are the more they are deflected.

  1. Detection

The ions that collide with the screen are detected

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24
Q

What are the conditions for Mass spectrometer

A
  1. The sample must be vaporized

2. The process must be done in a vacuum to prevent other ions interfering or atoms interfering

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25
Q

How to work out the relative atomic mass from mass spectrometry showing relative abundances

A

RAM= sum of ( isotopic mass x relative abundance) / Total relative abundance

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26
Q

How to work out the relative atomic mass from mass spectrometry showing relative percentage abundances

A

RAM = sum of (isotopic mass x percentage abundance)/ 100

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27
Q

What are some uses for mass spectrometry

A
  1. Drug testing to identify chemicals in the blood
  2. Pharmaceutical industry for quality control
  3. Planetary space probes to identify elements
28
Q

What are some rules surrounding electronic structure

A
  1. Electrons fill up the lowest energy subshells first ( s orbitals to d orbitals)
  2. The 4s subshell has a lower energy level than the 3D subshell- meaning it fills up first
29
Q

What are some exceptions to the electronic configuration rules

A
  1. Chromium and copper donate one of their 4s electrons to the 3D subshell as they are more stable with a full or half full 3D subshell
30
Q

What do Emission spectra provide evidence of?

A

Existence of quantum shells and that these shells are discrete

31
Q

What do successive ionisation energies provide evidence for

A

They provide evidence for the quantum shells within atoms and they suggest the group to which the element belongs to

32
Q

What do the first ionisation energies of succesive elements provide evidence for

A

They provide evidence for electron subshells

33
Q

What is an orbital

A

A region where an electron is most likely to be found

34
Q

What do the electrons do in orbitals

A

They spin in opposite directions sue to electron repulsion (spin-pairing)

35
Q

What is the shape of a s-orbital

A

Spherical shape

36
Q

What is the maximum number of electrons an S-orbital can hold and how many orbitals

A

1 orbital - 2 electrons

37
Q

What is the shape of a p-orbital

A

Dumbell shape

38
Q

What is the maximum number of electrons a P -orbital can hold and how many orbitals

A

3 orbitals - 6 electrons

39
Q

How many orbitals in a D orbital and what is the maximum number of electrons it can hold

A

5 orbitals - 10 electrons

40
Q

How many orbitals in an F orbital and how many electrons can it hold

A

7 orbitals- 14 electrons

41
Q

What is the order of the Em spectrum and the trend in wavelength and frequency

A
  1. Increasing in frequency

2. Radio- Microwaves- Infrared-visible light- UV- X-ray- Gamma ray

42
Q

What can electron shells also be known as and what is special about them

A

They can also be known as quantum shells and the levels are discrete

43
Q

What happens when an electron takes in energy in an emission spectra

A

When they take in energy they get excited and move to a higher energy level. When the electrons drop to a lower energy level A photon with a specific frequency is emitted which produces an emission spectra.

Different elements have different electron arrangements so the emission spectras are different

44
Q

What happens to the lines in emission spectra

A

They get closer together as frequency increases

45
Q

What are the 4 basic principles to Electron Shells

A
  1. Electrons can only exist in fixed orbits or shells
  2. Each shell has a fixed energy

3, When an electron moves between shells. EM radiation is emitted or absorbed

  1. Because the energy of shells is fixed, the radiation will have a fixed frequency
46
Q

How does Nuclear charge effect ionisation energy

A
  1. The more protons there are in the nucleus the more positively charged the nucleus is and the more attracted electrons are - higher ionisation energy
47
Q

How do the electron shells effect ionisation energy

A
  1. An electron closer to the nucleus has a stronger attraction to it - higher ionisation energy
48
Q

How does shielding effect ionisation energy

A

As the number of electrons between the outer electrons and the nucleus increases the outer electrons experience a lower attraction to the nucleus - lower ionisation energy

49
Q

Define second ionization energy

A

The energy required to remove 1 mole of electrons from 1 mole of gaseous 1+ ions to produce 1 mole of gaseous 2+ ions

50
Q

Describe the trend of first ionisation energies as you move down the group

A

The first ionisation energies decrease- meaning it is easier to remove outer electrons

51
Q

Why does the first ionisation energies decrease as you move down the group

A
  1. As you move down the group the atomic radius decreases due to increases shielding
  2. This means the outer electron are further away from the nucleus so the attraction to the nucleus is smaller- less energy is required to remove the outer electrons
52
Q

Why do successive ionisation energies increase within each shell

A

As electrons are being removed from an increasingly positive ion there will be less repulsion amongst the remaining electrons so they are held more strongly by the nucleus so more energy required to remove electrons

53
Q

How to know what group an element is from a list of successive ionisation energies or a graph

A
  1. From list = Find the biggest energy jump and pick the one it starts from
  2. From graph = Find how many electrons were removed before the first big jump this will tell you the group
54
Q

What does group number tell you

A

The number of electrons in the outer shell

55
Q

What does the period number tell you

A

The number of electron shells

56
Q

Describe the chemical properties across a group

A

Similar chemical properties as they have the same number of electrons int heir outer shell

57
Q

Describe what happens to atomic radius as you go across a period

A

Atomic radius decreases as you go across a period as the number of protons increases the positive charge of the nucleus increases. This means that electrons are more attracted to the nucleus and are pulled closer making the atomic radius smaller

58
Q

Describe what happens to Ionisation energy as you go across a period

A

As you go across a period the ionisation energy increases as it gets harder to remove electrons due to an increasingly positive nucleus as proton number increases

59
Q

Describe the drop in ionisation energy between Groups 2 and 3 due to subshell structure

A
  1. Describe their being an increase in nuclear charge between magnesium and aluminum.
  2. Aluminum’s outer electron sits in the 3p orbital rather than the 3s orbital so the electron is further away from the nuclues.
  3. The 3p orbital has additional shielding from the 3s2 electrons
  4. Both these factors override the increase in nuclear charge and the ionisation energy drops slightly
60
Q

Describe the drop in ionisation energy between Groups and 5 and 6 and use phosphorus and sulfur as examples

A
  1. Due to electron repulsion
  2. Phosphorus and Sulfur have identical amount of shielding however in Sulfur the outer electron is easier to remove due to 2 electrons occupying one 3p orbital.
  3. Due to this electron repulsion the electron is easier to remove and a lower ionisation energy is required compared to phosphorus
61
Q

What type of bonding occurs in metals and name the metals

A
  1. Metallic bonding

2. Li, Na, Be, Mg, Al

62
Q

Describe the trend in Melting and Boiling point

A
  1. The M.P and B.P increase across the period because the metallic bonds get stronger as the number of electrons increase so number of delocalised electrons increases.
  2. The atomic radius also decreases so there is a stronger attraction between metal ions and delocalised electrons.
63
Q

What are carbon and Silicon and describe their structures and B.P and M.P

A
  1. They are both Giant covalent lattice structures
  2. They have strong covalent bonds between their atoms and therefore a lot of energy is required to break all of these bonds. So their M.P and B.P are the highest in their periods.
64
Q

Describe the bonding in simple molecular structures and their M.P and B.P list some examples as well

A
  1. They have London forces
  2. Their melting and boiling points depend on the strength of their London forces. London forces are weak and are easily overcome so these elements have low melting and boiling points.
  3. N2, O2, F2, P4, S8 and Cl2
  4. More electrons in a molecule means stronger london forces so sulfur (s8) will have the strongest london forces
65
Q

Describe the the melting and boiling points of Noble gases

A
  1. The noble gases have the lowest M.P and B.P because they exist as individual atoms > very weak London Forces.

Chemistry (20 decks)

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