Atomic Structure And The Periodic Table

Question 1

What is an isotope?

Answer 1

Atoms which have the same atomic number but a different mass number

Isotopes contain the same number of protons but a different number of neutrons

Question 2

What is the atomic number of nitrogen?

Answer 2

7

Question 3

What is the atomic number of oxygen?

Answer 3

8

Question 4

What does the mass number represent?

Answer 4

The number of protons plus the number of neutrons

Question 5

What does the atomic number represent?

Answer 5

The number of protons in the nucleus of an atom

Question 6

How do you calculate the number of neutrons in an atom?

Answer 6

Mass number - Atomic number

Question 7

Fill in the blank: The mass numbers of isotopes of the same element will be _______.

Answer 7

different

Question 8

True or False: Every atom of an element has the same number of neutrons.

Answer 8

False

Question 9

What is the definition of relative isotopic mass?

Answer 9

The mass of an atom of an isotope of an element relative to one-twelfth of the mass of an atom of carbon-12

Question 10

What does relative atomic mass represent?

Answer 10

The weighted mean mass of an atom of an element compared to one-twelfth of the mass of an atom of carbon-12

Question 11

Fill in the blank: Relative isotopic mass is measured relative to _______.

Answer 11

one-twelfth of the mass of an atom of carbon-12

Question 12

Fill in the blank: Relative atomic mass is the _______ mean mass of an atom of an element.

Answer 12

weighted

Question 13

How many peaks will the mass spectrum of chlorine, Cl2, show?

Answer 13

5 peaks

Question 14

What charge do the species in the mass spectrum of chlorine, Cl2, have?

Answer 14

Positively charged

Question 15

Why are the species in the mass spectrum of chlorine, Cl2, positively charged?

Answer 15

Because it has lost an electron

Question 16

What are the principal energy levels labeled with?

Answer 16

Principal quantum number (n)

The lowest energy level is labeled as n=1, which is closest to the nucleus.

Question 17

What are the four types of sub-shells?

Answer 17

s, p, d, f

These sub-shells have different energy values and can hold varying numbers of electrons.

Question 18

What is the maximum number of electrons that the s sub-shell can hold?

Answer 18

2 electrons

The s sub-shell consists of one orbital.

Question 19

What is the maximum number of electrons that the p sub-shell can hold?

Answer 19

6 electrons

The p sub-shell consists of three orbitals.

Question 20

What is the maximum number of electrons that the d sub-shell can hold?

Answer 20

10 electrons

The d sub-shell consists of five orbitals.

Question 21

What is the maximum number of electrons that the f sub-shell can hold?

Answer 21

14 electrons

The f sub-shell consists of seven orbitals.

Question 22

Define an orbital.

Answer 22

A region in an atom that can hold up to two electrons with opposite spin

Orbitals are the specific regions within sub-shells.

Question 23

What is the shape of s orbitals?

Answer 23

S orbitals are spherical

S orbitals have a uniform shape around the nucleus.

Question 24

What is the shape of p orbitals?

Answer 24

P orbitals are dumb-bell shaped

P orbitals have two lobes extending in opposite directions.

Question 25

How many electrons can each orbital hold?

Answer 25

Each orbital can hold up to two electrons ## Footnote This applies to both s and p orbitals.

Question 26

What principle describes the order in which orbitals are filled?

Answer 26

The Aufbau principle ## Footnote This principle states that orbitals are filled in order of increasing energy.

Question 27

What is Hund's Rule?

Answer 27

Electrons prefer to occupy orbitals on their own, and only pair up when no empty orbitals of the same energy are available.

Question 28

What is the electron structure of Calcium (Ca)?

Answer 28

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁰ ## Footnote Calcium has a total of 20 electrons.

Question 29

What is the electron structure of Silicon (Si)?

Answer 29

1s² 2s² 2p⁶ 3s² 3p² ## Footnote Silicon has a total of 14 electrons.

Question 30

True or False: Electrons will always pair up in an orbital before occupying other orbitals of the same energy.

Answer 30

False

Question 31

What type of element is sodium based on its outer electrons?

Answer 31

s-block element ## Footnote Sodium's electron configuration is 1s² 2s² 2p⁶ 3s¹.

Question 32

Which block of the Periodic Table includes elements with outer electrons in a p subshell?

Answer 32

p Block ## Footnote Elements in this block have their outermost electrons in p orbitals.

Question 33

Which block of the Periodic Table includes elements with outer electrons in an s subshell?

Answer 33

s Block ## Footnote Elements in this block are characterized by having their outermost electrons in s orbitals.

Question 34

What block contains elements with outer electrons in a d subshell?

Answer 34

d Block ## Footnote These elements typically include transition metals.

Question 35

What is the definition of first ionisation energy?

Answer 35

The energy required to convert one mole of gaseous atoms into gaseous ions with a 1+ charge ## Footnote Example: Xie → X*i + e. Always include state symbols.

Question 36

What is the definition of second ionisation energy?

Answer 36

The energy required to convert one mole of 1+ gaseous ions into 2+ gaseous ions

Question 37

Is the process of ionisation exothermic or endothermic?

Answer 37

Endothermic

Question 38

Why do successive ionisation energies increase?

Answer 38

As electrons are removed, the resulting ion gets smaller, causing remaining outer electrons to experience greater nuclear attraction

Question 39

What factors affect ionisation energy?

Answer 39

* Atomic radius * Nuclear charge * Shielding

Question 40

How does atomic radius affect ionisation energy?

Answer 40

The larger the radius, the further the outer electrons are from the nucleus

Question 41

How does nuclear charge affect ionisation energy?

Answer 41

The higher the nuclear charge, the larger the attraction to the outer electrons

Question 42

What is the effect of shielding on ionisation energy?

Answer 42

Inner shells shield the outer electrons from the nucleus

Question 43

What does a big jump in ionisation energy indicate?

Answer 43

New shells

Question 44

What happens to the proton to electron ratio as electrons are removed?

Answer 44

It increases / effective nuclear charge increases

Question 45

What effect does an increased effective nuclear charge have on outer electrons?

Answer 45

Greater attraction between the nucleus and the outer electrons

Question 46

Fill in the blank: The inner shell electrons will always have the highest _______.

Answer 46

Ionisation Energy

Question 47

What is the trend in ionisation energy as electrons are removed?

Answer 47

General increase

Question 48

What is ionisation energy?

Answer 48

The energy required to remove an electron from an atom in its gaseous state ## Footnote Expressed in kJ mol-l

Question 49

What happens to ionisation energy as you move down a group?

Answer 49

Ionisation energy decreases ## Footnote Despite an increase in nuclear charge

Question 50

What factors contribute to the decrease in ionisation energy down a group?

Answer 50

* Outer electron is further from the nucleus * Outer electron is better shielded by inner shell electrons * Less force of attraction between outer electrons and nucleus ## Footnote This is due to increasing number of shells and shielding

Question 51

What happens to atomic radius and shielding as you move down a group?

Answer 51

Both atomic radius and shielding increase ## Footnote This results in decreased first ionisation energy

Question 52

What happens to ionisation energy as you move across a period?

Answer 52

Ionisation energy increases ## Footnote There is no increase in shielding and nuclear charge increases

Question 53

Why do noble gases have the highest first ionisation energy?

Answer 53

Their outer shells are full ## Footnote This results in the greatest force of attraction

Question 54

True or False: The nuclear charge decreases as you move down a group.

Answer 54

False ## Footnote The nuclear charge increases, but the effect of shielding and distance outweighs this

Question 55

What happens to the number of protons as you move across a period?

Answer 55

Increases ## Footnote This increase in protons affects various atomic properties.

Question 56

What happens to the atomic radius as the number of protons increases?

Answer 56

Decreases ## Footnote The increased nuclear charge pulls electrons closer to the nucleus.

Question 57

What generally happens to the first ionisation energy as you move across a period?

Answer 57

Increases overall ## Footnote This is due to increased nuclear charge and decreased atomic radius.

Question 58

What effect does increased nuclear charge have on ionisation energy?

Answer 58

Increases ionisation energy ## Footnote A higher nuclear charge means a stronger attraction between the nucleus and electrons.

Question 59

What is the reason for the decrease in ionisation energy between Be and B?

Answer 59

Easier to remove an electron from the 2p orbital than from the 2s orbital ## Footnote This is due to electron configurations and energy levels.

Question 60

Which orbital does the electron removed from boron come from?

Answer 60

2p orbital ## Footnote Electrons in the 2p orbital are at a higher energy level than those in the 2s orbital.

Question 61

Why does nitrogen have a higher ionisation energy than oxygen?

Answer 61

Nitrogen has a half-filled 2p sub-shell ## Footnote A half-filled subshell is more stable than one with paired electrons.

Question 62

What is a key characteristic of a half-filled subshell?

Answer 62

More stable ## Footnote This stability affects the ionisation energy of elements.

Question 63

Fill in the blank: The atomic radius ______ as the number of protons increases across a period.

Answer 63

decreases

Question 64

True or False: Increased nuclear charge leads to decreased ionisation energy.

Answer 64

False ## Footnote Increased nuclear charge actually leads to increased ionisation energy.