Chapter 3: Periodicity

Question 1

What are element in the periodic table arranged by?

Answer 1

Increasing atomic number/number of protons

Question 2

What is the relative atomic mass of an element?

Answer 2

The weighted average mass of an atom in comparison to Carbon 12.

Question 3

What does the group/period of an element tell us?

Answer 3

The group numbers shows the number of valence electrons in the outer shell whereas the period number represents the number of main energy levels each atom contains.

Question 4

What are metalloids and where are they located on the periodic table?

Answer 4

Metalloids have properties that include those of metals an non-metals are form a dividing staircase between the metals and non-metals

Question 5

What are 5 group names in the periodic table?

Answer 5

Group 1: Alkali metals
Group 2: Alkali Earth metals
Group 17: Halogens
Group 18: Noble Gases
Group 3-12: Transition metals (excluding zinc)
Lanthanoids and Actinoids

Question 6

What are some physical properties of metals?

Answer 6

Solids at room temperature
Malleable
Metallic lustre
High electrical conductivity
High densities
Low ionization energies
Low electronegativity values

Question 7

What are some physical properties of non-metals?

Answer 7

Mostly gases at room temp
Dull
Brittle
Low densities
Poor electrical conductivity
High ionization energies
high electronegative values

Question 8

What are some chemical properties of metals?

Answer 8

Lose electrons to form positive cations
Behave as reducing agents
Form ionic bonds wihj non-metals
Form basic oxides

Question 9

What are some chemical properties of non-metals?

Answer 9

Gain electrons to form negative anions
Behave as oxidizing agents
Form covalent bonds with other non-metals
For acidic oxides

Question 10

What are some physical properties of metalloids?

Answer 10

Solids at room temp
Some have metallic lustre
Brittle
Intermediate electrical conductivity
Moderate density
Intermediate ionization energies + electronegativity values

Question 11

What are some chemical properties of metalloids?

Answer 11

Produce amphoteric or weakly acidic oxides

Question 12

What is electron shielding?

Answer 12

Electron shielding refers to the effect that inner energy level electrons have on the outer valence electrons, shielding them from the full attraction of the nucleus.
This valence electron requires less energy to remove than the inner electron (ionization energy)

Question 13

What is the trend in electron shielding in the periodic table?

Answer 13

Electronic shielding remains constant across a period due to the elements consisting of the same number of main energy levels, thus effectively experiencing similar shielding from the inner shells. Whereas down a group, electron shielding increases as the number of main energy levels increase.

Question 14

What is effective nuclear charge (Z) ?

Answer 14

The net positive charge experienced by valence electrons
= Atomic number - number of shielding electrons

Question 15

What is the trend in effective nuclear charge across the periodic table?

Answer 15

Increases across a period as atoms have the same number of main energy levels but increasing protons, thus resulting in increasing ionization energies across a period.
Down a group, it remains the same as the number of main energy levels increase along with increasing protons. The ionization energies increase however due to increasing atomic radii.

Question 16

What are the trends in atomic radii across the periodic table?

Answer 16

Atomic radii decrease across a period as effective nuclear charge increases and thus the attraction between the nucleus and valence electrons increase.
Atomic radii increases as the number of main energy levels increase, thus resulting in valence electrons being less attracted to the nucleus.

Question 17

What are the trends in ionic radii down the periodic table?

Answer 17

Increase in ionic radii down a group, due to the increase in occupied main energy levels.

Question 18

What are the trends in ionic radii across the periodic table?

Answer 18

Decrease in ionic radii across a period, due to the increase in attraction between the nucleus and electrons.

Question 19

What are ions that have the same electronic structure called?

Answer 19

An isoelectronic species

Question 20

What are the trends in size of parent atoms and their positive ions? (metals)

Answer 20

The parent atoms are larger than the ion itself

Question 21

What are the trends in size of parent atoms and their negative ions? (non-metals)

Answer 21

Negative ions are larger than their parent ion themselves due to having more electrons, therefore there are weaker forces of electrostatic attraction between the electrons and nucleus.

Question 22

What is the first ionization energy?

Answer 22

The energy required to remove 1 mole of electrons from one mole of gaseous atoms to produce 1 mol of gaseous 1+ ions.

Question 23

Why are the values positive for first ionization energies?

Answer 23

As the energy must be supplied in order to overcome electrostatic attraction between the nucleus and outer electrons.

Question 24

What is the trend in ionization energies in the periodic table?

Answer 24

Ionization energies increase across a period, with the noble gases having the greatest ionization energies as the nuclear charge increases and the atomic radii decreases.
However decreases down a group as the number of main energy levels increase.

Question 25

What are the exceptions to the trends in ionisation energy and why?

Answer 25

There is a slight dip in ionisation energy across a period, as while elements consist of the same main energy levels, they may have orbitals that are slightly higher in energy, thus making it easier to remove electrons.
IE Be = 1s2 2s2 & B = 1s2 2s2 2p1, therefore B has a lower ionisation energy than Be.

Question 26

Why does oxygen have a lower ionisation energy than nitrogen?

Answer 26

N = 1s2 2s2 2p3
O = 1s2 2s2 2p4
Therefore an electron in a doubly occupied orbital experiences repulsion from the other electron and thus is easier to remove than an electron in a half-filled orbital.

Question 27

What are the trends in electronegativity in the periodic table?

Answer 27

Electronegativity increases across a period due to increasing nuclear charge, resulting in stronger attraction between the nucleus and bonding electrons.
Electronegativity decreases down a period due to the increase in atomic radii and main energy levels, therefore weaker forces of electrostatic attraction.

Question 28

What is electronegativity?

Answer 28

A measure of the attraction an atom has for a bonding pair of electrons.

Question 29

What are the trends in melting point in the periodic table?

Answer 29

Melting point depends on type of bonding.
Molecular covalent < metallic bonding < giant covalent.

Question 30

What are the trends in metallic character in the periodic table?

Answer 30

Down a group, increasing metallic character.
Across a period, decreasing metallic character.

Question 31

What is metallic character?

Answer 31

How easily an atom loses electrons to form an ion.

Question 32

What is the 1st electron affinity?

Answer 32

Energy released when one mole of electrons is added to one mole of gaseous atoms to form one mole of 1- ions.
It is exothermic therefore values are negative.
Non-metals tend to have greater exothermic values due to their tendency to form negative ions.

Question 33

What are the trends in 1st electron affinity in the periodic table?

Answer 33

Across a period, electron affinity increases as atomic radii decreases and electron shielding is relatively constant.
Down a group, electron affinity decreases as atomic radii decreases and electron shielding increases, therefore less energy is released.

Question 34

Why are 2nd electron affinity values positive?

Answer 34

As energy must be supplied to overcome forces of repulsion due to adding an electron to an already negative ion.

Question 35

Why does Fluorine have a lower electron affinity than Chlorine?

Answer 35

As it has a very small atomic radius, therefore experiences more repulsion when an electron is added

Question 36

What are some physical properties of Group 1 metals?

Answer 36

Soft shiny metals that can be easily cut.
Decreasing melting point down the group due to weaker metallic bonding.
Atomic and ionic radii increase due to increase in number of main energy levels.
Low densities, can float on water.

Question 37

What are some chemical properties of Group 1 metals?

Answer 37

Reactivity increases down the group.
React with water to produce alkaline solutions.
Ionisation energies decrease.
Metallic character increases and electronegativity decreases.

Question 38

What is produced when a group 1 metal reacts with a halogen?

Answer 38

Salts are produced where ionic bonds are formed between the oppositely charged ions.
I.E sodium chloride

Question 39

What are some physical properties of Group 17 halogens?

Answer 39

Melting + Boling points increase down the group due to increasing molar mass, stronger London dispersion forces.
Atomic + ionic radii increases, electronegativity and ionisation energy decreases.
Electron affinity becomes less exothermic.

Question 40

What are some chemical properties of Group 17 halogens?

Answer 40

Reactivity decreases and elements at the top are stronger oxidizing agents.
A more reactive halogen will displace a less reactive one from a solution.

Question 41

What are the changes in bonding across the periodic table?

Answer 41

Bonding changes from ionic to covalent to molecular covalent due to decreasing electronegativity.

Question 42

What are the trends in pH of the oxides across the periodic table?

Answer 42

Basic —> Amphoteric —-> Acidic