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Chemistry > Atomic Structure and the Periodic Table > Flashcards

Atomic Structure and the Periodic Table Flashcards

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1
Q

What is atomic number (Z)

A

The number of protons in the nucleus of an atom of an element

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2
Q

What is the mass number of an element

A

is the sum of the number of protons and the number of neutrons in the nucleus of an atom

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3
Q

What is an isotope

A

Are atoms of the same element witht the same number of protons but a different mass number due to a different number of protons

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4
Q

What is the symbol,

relative mass/charge

and position in atoms of a

Proton

A

p

1

+1

nucleus

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5
Q

What is the symbol,

relative mass/charge

and position in atoms of a

Neutrons

A

n

1

0

nucleus

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6
Q

What is the symbol,

relative mass/charge

and position in atoms of a

Electron

A

e-

1/1840

-1

energy levels surrounding the nucleus

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7
Q

Relative atomic mass (Ar)

A

The weighted mean mass of an atom of an element compared to 1/12 of the mass of an atom of carbon-12

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8
Q

Relative isotopic mass

A

The mass of an individual atom of a particular isotope relative to 1/12 of the mass of an atom of carbon-12

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9
Q

What does a mass spectrometer do

How does it work

A

measures the mases of atoms and molecules

Produces positive ions, which are deflected by a magnetic feild, according to the m/z ratio

It also calculates relative abundance in %

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10
Q

Why are mass spectrometers important in terms of isotopes

A

determine the exact values of relative masses of isotopes

As well as the % abundance

This information can be used to calculate relative atomic mass

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11
Q

How can the relative molecular mass of diatomic molecules be measured

A

In a mass spectrometer

you can obtain Mr of element/compound by observing peaks with largest m/z ratio

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12
Q

What is the molecular ion peak

A

the peak with the highest m/z ratio in the mass spectrum

M peak

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13
Q

How would you work out the number of neutrons from atomic (z) and mass numbers

A

mass number - atomic number = neutrons

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14
Q

How can you work out mass number?

A

atomic (protons z) + neutrons = mass number

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15
Q

What would the mass spectrometry look like for chlorine and why?

A

3:1 ratio for Chlorine 35 and 37

Then a 9:6:1 ratio at 70,72,74 due to Chlorine being diatomic.

(35+35=70) (35+37=72) (37+37=74)

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16
Q

Why is there a 9:6:1 ratio with peaks at 70, 72 and 74 on the chlorine mass spectrometry

A

It is a 3:4 ratio for Chlorine 35 and 37

(35+35) 70: 3/4 x 3/4 = 9/16

(35+37) 72: 3/4 x 1/4 = 3/16

(37+37) 74: 1/4 x 1/4 = 1/16

9:6:1

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17
Q

What would the mass spectrometry look like for bromine and why

(two isotopes are 79 and 81)

A

two peaks at 79 and 81 at equal heights at a 1:1 ratio

(79+79) = 158: 1/2 x 1/2 = 1/4

(79+81) = 160: 2(1/2 x 1/2) = 2/4

(81+81) = 162: 1/2 x 1/2 = 1/4

1:2:1

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18
Q

How do you work out relative atomic mass from relative isotopic mass and percentage abundance on a mass spectrometer

A

(relative isotopic mass x percentage abundace) / 100

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19
Q

Work out the relative atomic mass of lithium when 7.59% of the sample has 6.015 isotopic mass and 92.41% has a 7.016 isotopic mass

A

(6.015 x 7.59) + (7.016 x 92.41) / 100

= 6.94

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20
Q

What is the molecular ion peak

A

is the peak with the highest m/z ratio on the spectrum (M)

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21
Q

How would you work out the relative molecular mass of Chlorine, when you have 9:6:1 ratio of peaks at 70,72,42

A

(9 x 70) + (6 x 72) + (1 x 74) / 16

= 71

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22
Q

What is a quantum shell

A

defines the energy level of an electron

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23
Q

What is an orbital

A

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

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24
Q

What is electron configuration

A

shows the number of electrons in each sublevel in each energy level of an atom

25
Q

What is Hund’s rule

A

states that electrons will occupy the orbitals singly before pairing takes place

26
Q

What is the Pauli Exclusion Principle

A

states that two electrons cannot occupy the same orbital unless they have opposite spins. Electrons spins are shown by single-headed arrows in opposite directions

27
Q

What are the 4 quantum shell and how many electrons can go in each

A

S = 2 (1 x 2)

P = 6 (3 x 2)

D = 10 (5 x 2)

F = 14 (7 x 2)

x 2 because each orbital can hold 2 electrons

28
Q

What is the shape of the S orbital

A
29
Q

What is the shape of the P orbital

A
30
Q

How many electrons can fill the first 4 quantum shells

A

1st : 2 (1s²)

2nd : 8 (2s² 2p6)

3rd: 18 (3s² 3p6 3d10)
4th: 32 (4s² 4p6 4d10 4f14)

31
Q

What is the exceptions to the allocation of electrons in the electron configuration

A

the 4s has a lower energy level than the 3d shell therefore the 4s shell fills before the 3d shell does

32
Q

How do you configure electrons in boxes

A

an electron will fill an orbital singly before pairing takes place

only two electrons can occupy and orbital with opposite spins

33
Q

What is the first ionisation energy

A

is the energy required to remove an electron from each atom in one mole of atoms in gaseous state

34
Q

What is second ionisation energies

A

is the energy required to remove an electron from each singly charged positive ion in one mole of positive ions in gaseous state

35
Q

How can we signify first ionisation energy of element A

A

A(g) + A+(g) + e-

36
Q

How can we signify second ionisation energy

A

A+(g) → A2+(g) + e-

37
Q

Where did evidence for different energy levels initially come from

A

atomic emission spectra

Heating atoms in gaseous state, moving electrons to a higher energy level. When they return back they release EM radiation.

There are many specific frequencies produced per element, indicating it requires a different amount of energy to remove electrons

38
Q

How can you tell what group an element is in by looking at its ionisation energies

A

big jumps in energies is a piece of evidence for quantum shells

39
Q

How can you tell where groups occur on a logarithm for successive ionisation for Sodium

A

quantum shell configuration: 2, 8, 1

Jump from 1st to 2nd, due to the outer electron being removed from the highest energy level of the third quantum

Jump from 9th to 10th, due to movement from the 2nd energy level to the first

40
Q

Why do successive ionisation energies increase in magnitude

A

The first electron lost during ionisation is furthest from the influence of the nucleus

Out electrons also exits in high energy levels, so will require little to remove it

41
Q

What is the equation for working out ionisation energy (IE)

A

IE = energy of electron when removed - energy of electron when in orbital

42
Q

When does electron-electron repulsion occur

A
  • two electrons in the same orbital
  • electrons in different orbitals in a quantum shell
  • Electrons in adjacent quantum shells
43
Q

Which first ionisation energy will be the largest in Hydrogen or Helium

Hydrogen(1s1)

Helium (1s2)

A

Outer electrons are in the same orbital

However, helium contains 2 electrons, increasing electron-electron repulsion - each electron is said to shield one another from nuclear charge

Nuclear charge of helium is double that of hydrogen

Effect of nuclear charge us greater than shielding, so helium has a higher ionisation energy

44
Q

Which ionisation energy is greater, Helium or lithium

Helium (1s2)

Lithium (1s2 2s1)

A

Lithium has a greater nuclear charge

Lithium’s outer electron is in the 2s orbital, being in the second quantum shell, having a higher energy level. Therefore lower ionisation energy

Will also experience electron-electron repulsion from the 1s orbital

So Helium has a higher ionisation energy

45
Q

What effect the energy the electron has, hence the first ionisation energy

A
  • Orbital the electron exists within
  • Nuclear charge of the atom (no of protons)
  • Repulsion (shielding) experienced by electrons from all other electrons present
46
Q

What is the trend in ionisation energies across a period and why

From (Li to Ne) period 2

A

Nuclear charge will increase, as a proton will be added each time

However, one more electron is also added to same quantum shell, increasing electron-electron repulsion

Increase in nuclear charge is greater than that of electron repulsion, so first ionisation energy increases across period 2

47
Q

Trend in ionsation energy down a group and why

(Li to Cs) group 1

A

Nuclear charge increase, as number of protons increases meaning increased attraction with nucleus - increase in ionisation energies

However, new quantum shell is added, each occasion, which has a higher energy level and increases electron-electron repulsion between shells

Therefore, first ionisation energy decreases down the group

48
Q

Why do elements within groups have similar chemical properties

A

same outer electron confirguration

49
Q

All elements in a main group, contain the same electron configuration

If n = the period

Write the electron configuration for groups 1-8

A

Group 1 : ns1

Group 2 : ns2

Group 3 : ns2 np1

Group 4 : ns2 np2

Group 5 : ns2 np3

Group 6 : ns2 np4

Group 7 : ns2 np5

Group 8 : ns2 np6

50
Q

Which groups of the periodic table does the S block contain

A

Groups 1 and 2

51
Q

Which groups of the periodic table does P-block contain

A

Groups 3-8

52
Q

What elements are contained within d-block

A

Transition metals

Period 4 (Sc to Zn)

Period 5 (Y to Cd)

53
Q

Defintion of peridicitiy

A

is a regularly repeating pattern of atomic, physical and chemical properties with increasing atomic number

54
Q

What is the trend across the periodic table for size of nucleus

A

Size of radius decrease across each period

This is because the number of protons increases, increasing attractive forces between outer electrons

Although is offset through electron-electron repulsion as an electron is added each time

55
Q

What is the trend in melting temperatures across period two

A

Increase from Li to C

Due to molecules being more charged

Decrease from N2 to F2 because they exist as simple molecular diatomic molecules

56
Q

What is the trend in boiling temperature across group 2

A

Increase from Li to C

Due to more charged molecules and moving from metallic to covalent bonding

Decrease from N2 to F2 due to simple molecular diatomic covalent bonding

57
Q

What is the trend with boiling and melting temperatures across period 3

A

Increase from Na to Si

Due to increase in charge and moving from metallic giant lattices to covalent giant lattices

Then a decrease from P2 to Cl2 due to simple molecular covalent bonding

58
Q

Describe the graph for first ionisation energies for the first 3 periods

A

First ionsation energies highest for Nobel Gases (He, Ne, Ar)

Lowest for group one elements (Li and Na)

Anomaly increase at Helium

Group 2 (ns2) and group 3 (ns2 np1) the np1 in group 3 increases increases electron-electron repulsion therefore group 3 has lower ionisation energies than group 2

59
Q

Why do nitrogen and oxygen have an anomaly when calculating ionisation energies

A

N: 1s2 2s2 2p3

O: 1s2 2s2 2p4

First electron removed from Oxygen is paired, which has increased electron-electron repulsion

Less energy is required to remove the first electron because it is the same quantum shell and energy level

Chemistry (12 decks)

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