Atomic Structure and the Periodic Table

Question 1

What is atomic number (Z)

Answer 1

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

Question 2

What is the mass number of an element

Answer 2

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

Question 3

What is an isotope

Answer 3

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

Question 4

What is the symbol,

relative mass/charge

and position in atoms of a

Proton

Answer 4

p

1

+1

nucleus

Question 5

What is the symbol,

relative mass/charge

and position in atoms of a

Neutrons

Answer 5

n

1

0

nucleus

Question 6

What is the symbol,

relative mass/charge

and position in atoms of a

Electron

Answer 6

e^-

1/1840

-1

energy levels surrounding the nucleus

Question 7

Relative atomic mass (A_r)

Answer 7

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

Question 8

Relative isotopic mass

Answer 8

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

Question 9

What does a mass spectrometer do

How does it work

Answer 9

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 %

Question 10

Why are mass spectrometers important in terms of isotopes

Answer 10

determine the exact values of relative masses of isotopes

As well as the % abundance

This information can be used to calculate relative atomic mass

Question 11

How can the relative molecular mass of diatomic molecules be measured

Answer 11

In a mass spectrometer

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

Question 12

What is the molecular ion peak

Answer 12

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

M peak

Question 13

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

Answer 13

mass number - atomic number = neutrons

Question 14

How can you work out mass number?

Answer 14

atomic (protons z) + neutrons = mass number

Question 15

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

Answer 15

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)

Question 16

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

Answer 16

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

Question 17

What would the mass spectrometry look like for bromine and why

(two isotopes are 79 and 81)

Answer 17

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

Question 18

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

Answer 18

(relative isotopic mass x percentage abundace) / 100

Question 19

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

Answer 19

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

= 6.94

Question 20

What is the molecular ion peak

Answer 20

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

Question 21

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

Answer 21

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

= 71

Question 22

What is a quantum shell

Answer 22

defines the energy level of an electron

Question 23

What is an orbital

Answer 23

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

Question 24

What is electron configuration

Answer 24

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

Question 25

What is Hund's rule

Answer 25

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

Question 26

What is the Pauli Exclusion Principle

Answer 26

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

Question 27

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

Answer 27

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

Question 28

What is the shape of the S orbital

Answer 28

Question 29

What is the shape of the P orbital

Answer 29

Question 30

How many electrons can fill the first 4 quantum shells

Answer 30

1st : 2 (1s²) 2nd : 8 (2s² 2p⁶) 3rd: 18 (3s² 3p⁶ 3d¹⁰) 4th: 32 (4s² 4p⁶ 4d¹⁰ 4f¹⁴)

Question 31

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

Answer 31

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

Question 32

How do you configure electrons in boxes

Answer 32

an electron will fill an orbital singly before pairing takes place only two electrons can occupy and orbital with opposite spins

Question 33

What is the first ionisation energy

Answer 33

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

Question 34

What is second ionisation energies

Answer 34

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

Question 35

How can we signify first ionisation energy of element A

Answer 35

A_(g) + A⁺_(g) + e^-

Question 36

How can we signify second ionisation energy

Answer 36

A⁺_(g) → A²⁺_(g) + e^-

Question 37

Where did evidence for different energy levels initially come from

Answer 37

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

Question 38

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

Answer 38

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

Question 39

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

Answer 39

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

Question 40

Why do successive ionisation energies increase in magnitude

Answer 40

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

Question 41

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

Answer 41

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

Question 42

When does electron-electron repulsion occur

Answer 42

* two electrons in the same orbital * electrons in different orbitals in a quantum shell * Electrons in adjacent quantum shells

Question 43

Which first ionisation energy will be the largest in Hydrogen or Helium Hydrogen(1s¹) Helium (1s²)

Answer 43

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

Question 44

Which ionisation energy is greater, Helium or lithium Helium (1s²) Lithium (1s² 2s¹)

Answer 44

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

Question 45

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

Answer 45

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

Question 46

What is the trend in ionisation energies across a period and why From (Li to Ne) period 2

Answer 46

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

Question 47

Trend in ionsation energy down a group and why (Li to Cs) group 1

Answer 47

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

Question 48

Why do elements within groups have similar chemical properties

Answer 48

same outer electron confirguration

Question 49

All elements in a main group, contain the same electron configuration If n = the period Write the electron configuration for groups 1-8

Answer 49

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

Question 50

Which groups of the periodic table does the S block contain

Answer 50

Groups 1 and 2

Question 51

Which groups of the periodic table does P-block contain

Answer 51

Groups 3-8

Question 52

What elements are contained within d-block

Answer 52

Transition metals Period 4 (Sc to Zn) Period 5 (Y to Cd)

Question 53

Defintion of peridicitiy

Answer 53

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

Question 54

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

Answer 54

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

Question 55

What is the trend in melting temperatures across period two

Answer 55

Increase from Li to C Due to molecules being more charged Decrease from N₂ to F₂ because they exist as simple molecular diatomic molecules

Question 56

What is the trend in boiling temperature across group 2

Answer 56

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

Question 57

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

Answer 57

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

Question 58

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

Answer 58

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

Question 59

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

Answer 59

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