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AEC > Atomic Structure & The Periodic Table > Flashcards

Atomic Structure & The Periodic Table Flashcards

1
Q

Element

A

A substance that cannot be broken down into two or more different substances

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2
Q

Atom

A

Smallest uncharged particle of an element

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3
Q

Compound

A

Made of two or more elements chemically joined together

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4
Q

Ion

A

Charged atom or group of Atoms

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5
Q

What is the structure of an atom?

A

It has a small, central , positively charged nucleus (protons + neutrons) with electrons orbiting.

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6
Q

What is the relative mass of a proton ?

A

1

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7
Q

What is the relative mass of a neutron?

A

1

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8
Q

What is the relative mass of an electron?

A

1/1840

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9
Q

What is the relative charge of protons?

A

+1

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10
Q

What is the relative charge of neutrons?

A

0

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11
Q

What is the relative charge of electrons?

A

-1

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12
Q

What is meant by the term atomic number?

A

Is the number of protons in the nucleus of an atom of that element

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13
Q

What is meant by the term mass number?

A

The sum of the number of protons and neutrons in a nucleus of an atom

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14
Q

How do we determine the number of each sub atomic particle from atomic number or mass number?

A

Atomic Number = Number of Electrons

Mass Number = Neutrons + Protons

Number of Neutrons = Mass Number - Atomic Number

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15
Q

Define the term Isotope

A

Are atoms of the same element that have the same number of protons but different number of neutrons

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16
Q

Define relative isotopic mass based on the 12c scale

A

The mass of an atom of that isotope divided by 1/12 the mass of a carbon-12 atom

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17
Q

Define relative atomic mass (Ar)based on the 12c scale

A

The weighted average mass of an atom of that element divided by 1/12 the mass of a carbon-12 atom

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18
Q

Define relative molecular mass (Mr)?

A

Of an element or a compound is the average mass of a molecule of that element or compound divide by 1/12 the mass of a carbon-12 atom

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19
Q

When should the term Relative formula mass be referred to ?

A

Used for compounds with giant structures

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20
Q

What does the letter Z represents?

A

The atomic number of an atom

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21
Q

What does the letter A represents?

A

Mass Number

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22
Q

What is the relative atomic mass quoted on the periodic table table showing ?

A

The weighted average of all the isotopes.

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23
Q

For each isotope what can the mass spectrometer measure?

A

m/z value + abundance

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24
Q

How do you calculating the relative atomic mass from relative abundance of isotopes from mass spectrometry

A

R.A.M = Σ ( isotopic mass x % abundance) / 100

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25
Q

If two electrons are removed from a particle forming a 2+ ion… what would the m/z value be?

A

Half the mass

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26
Q

Predicting mass spectra for diatomic molecules like chlorine

A
  1. Identify the peaks
  2. The higher m/z peaks corresponds to the addition of the molecules
  3. Apply the RAM formula divide by 2
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27
Q

How can mass spectrometry be used to determine the relative molecular mass of a molecule ?

A

The molecules will break up and give a series of peaks caused by the fragments.

The peak with the largest m/z value is the complete molecule → molecular ion

28
Q

What leads to the m+1 peak?

A

Existence of the small percentage of carbon-13

29
Q

Define first ionisation energy

A

The energy required when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge

30
Q

Define successive ionisation

A

number of electrons removed from the successive no of shells and the energy involved

31
Q

What are the factors that affect ionisation energy?

A
  • Attraction of the nucleus → more protons , greater attraction
  • Distance of electrons from the nucleus → bigger atom, further outer electron, weak attraction
  • Shielding → electron in outer shell repelled by electrons in complete inner shells
32
Q

How are ionisation energies influenced by number of protons? (Nuclear charge)

A
  • The greater the number of protons, the greater the attraction of the electrons to the nucleus
  • The greater attraction of the electron → more energy is needed to remove electron
  • Hence IE is greater
33
Q

How are ionisations energies influenced by the number of electron shielding? (Shielding)

A
  • The greater the number of electrons between the nucleus + outer shell → lower the effective nucleus charge
  • Because the positive charge felt by the electron is reduced by the electrons between them
  • This means that the greater the number of electrons → lower the IE
34
Q

How are ionisation energies influenced by the number of electron sub shells From which electron is removed? (Distance from nucleus)

A
  • Electrostatic attraction decreases sharply with distance
  • Hence less energy needed to remove electrons which are further away
  • So as distances increases, IE decreases
35
Q

What are the reasons for the general increase in first ionisation energies across a period?

A

As one goes across a period, the number of protons increases

So the effective attraction of the nucleus is greater

The electrons are being added to the same shells

Which has the same shielding effect

And the electrons are pulled in closer to the nucleus

36
Q

What are the reasons for the decrease in first ionisation energy down a group?

A

As one goes down the group

The outer electrons are found in shells

further from the nucleus

and are more shielded

so the attraction of the nucleus becomes smaller

37
Q

What are the 3 ways in which the ideas about electronic configuration developed from?

A
  • Atomic emission spectra - existence of quantum shells
  • Successive IE - quantum shells + group to which elements belong to
  • 1st IE of Successive elements - electron sub shells
38
Q

How did the atomic emission spectra provide evidence for the existence of quantum shells?

A

Presence of certain lines → an electron can only adopt certain discrete energy levels

→ hence quantum shells (Defines energy lvl of electrons)

39
Q

How did successive ionisation energies provide evidence for the existence of quantum shells and the group to which the elements belongs?

A

Evidence for number of electrons in outer orbit

→ Big jumps between 2 IEs show how many electrons in their outer shell ,

→ the difference shows that it is harder to remove than the preceding hence moving orbits

40
Q

Why is using successive IE not a viable method for d-block elements?

A

Because they all have 2 electrons in their outer orbit.

41
Q

Which elements are exceptions to the d-block ?

A

Chromium + Copper which have only one outer shell

42
Q

How did the first ionisation energy of successive elements provide evidence for electron sub-shells? (Check answer again)

A

The bigger jump to the number of electrons being removed → electron is in a inner shell closer to the nucleus

→ more attracted to nucleus than the previous electron being removed

43
Q

What are the number of electrons that can fail up the first four quantum shells?

A

1st shell - 2

2nd shell- 8. (2+6)

3rd shell - 18 (2+6+10)

4th shell - 32 (2+6+10+14)

44
Q

What is an orbital?

A

A region within an atom that can hold up to two electrons with opposite spins

45
Q

What is the shape of an S-orbital?

A

Spherical

With 2 electrons moving anywhere within the sphere

46
Q

What is the shape of an P-orbital?

A

3 orbitals 90degrees to each other

Dumbbells shape

Electrons can move anywhere

47
Q

How many electrons occupy the s sub shell?

A

2 electrons

48
Q

How many electrons occupy the p sub shell?

A

6 electrons

49
Q

How many electrons occupy the d sub shell?

A

10 electrons

50
Q

How do electrons fill sub shells?

A

They fill them singly - one each

51
Q

What must two electrons in the same orbital have?

A

Before pairing up , the two electrons in the same orbital must have opposite spins

52
Q

How do we predict the electron configuration?

A

Using 1s notation + electron in box

53
Q

How can elements be classified?

A

Into s, p and d-block elements

54
Q

What does the electron configuration determine?

A

The chemical properties of an element

55
Q

What is periodicity?

A

A repeating pattern of physical or chemical properties going across a period

56
Q

What is the MP+BP of element in period 2+3?

A

The melting temperatures of the elements fist four elements then fall dramatically at fifth

57
Q

Why is the MP+BP of period 2 +3 elements like this?

A

The melting temperatures of the metals increase across a group as more electrons are involved in bonding. The non-metals - simple molecular elements Consist of weak London forces between the molecules hence litre less energy needed to break them.

58
Q

What is the trend in ionisation energy for period 2+3 elements?

A

The general trend across is to increase across a period

59
Q

Why is the trend for period 2 + 3 elements like this ?

A

This is due to the increasing number of protons as the electrons are being added to the same shell

60
Q

Why is there a small drop between Mg + Al in IE?

A

Mg has a its outer shell electrons in the 3s sub shell , whereas Al is starting to fill the 3p sub shell. Al’s electrons is slightly easier to remove because the 3p electrons are higher in energy.

61
Q

Why is there a small drop between Phosphorus + Sulfer’s in IE?

A

Sulfurs’s outer electron is being paired up with another electron in the same 3p orbital.

When the second electron is added to an orbital there is a slight repulsion between the two negatively charged electrons which makes the second electron easier to remove.

62
Q

How do we illustrate periodicity In terms of Electron Configuration?

A

Elements in each period have the same valence shell

63
Q

How do we illustrate periodicity In terms of Atomic Radii?

A

As the atomic number increase the atomic radii decreases as the positive charge will bring electrons closer to the nucleus

64
Q

How do we illustrate periodicity In terms of MP+BP?

A

Across period valency increase so the metals atoms can delocalise more electrons to form more positively charged cations and a bigger sea of delocalised electrons. ( data )

65
Q

How do we illustrate periodicity In terms of First ionisation energy?

A

Using the date - increase

AEC (19 decks)

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