Chemistry Module 3 Trends in Period 3

Question 1

What are the 7 trends across the period?

Answer 1

-Atomic Radius
-Electronegativity
-First Ionization Energy
-Structure and Bonding
-Melting Point
-Conductivity
-Density

Question 2

What are the three main facts about period 3 elements?

Answer 2

-They have three shells
-The number of protons increases across the period
-The number of electrons increases across the period

Question 3

What is the definition of the term Atomic Radius?

Answer 3

This is the distance between the nuclei and the valence electron of an atom of an element.

Question 4

In the periodic table moving from left to right across the period, what impact is played on the atomic radius?

Answer 4

The Atomic Radius decreases from left to right because:
-The nuclear charge increases as the number of protons increases
-There is no change in shielding and the number of shells remains the same
-There is an increased attraction between the nucleus and the electrons

Question 5

What is electronegativity?

Answer 5

The tendency for an atom to attract a bonding pair of electrons.

Question 6

When moving from left to right in the period what is the effect on electronegativity?

Answer 6

Electronegativity increases because:
-There is a decrease in the atomic radius.
-There is an increase in nuclear
These factors make it easier to pull in or attract electrons

Question 7

What is the first ionization energy?

Answer 7

This is the energy required to remove the outermost / highest energy electron from the valence shell of an atom.

Question 8

As we move from left to right in the period what happens to the first ionization energy?

Answer 8

The first ionization energy increases because:
-The Atomic Radius decreases
- The force of attraction between the nucleus and the valence electron increases (nuclear charge)
Due to these factors, more energy is needed to remove an electron from the valence shell.

Question 9

What are the 2 anomalies for the first ionization energy?

Answer 9

They are Aluminium and Sulphur

Question 10

What are the electronic configurations of Mg and Al?

Answer 10

Mg: 1s2, 2s2, 2p6, 3s2
Al: 1s2, 2s2 2p6, 3s2 3p1

Question 11

Why is there a drop between Mg and Al?

Answer 11

The outer electron moved to the 3p orbital. In this orbital, it takes less energy to remove an electron than the 3s orbital.

Question 12

What is the electronic configuration of P and S?

Answer 12

P: 1s2, 2s2 2p6, 3s2 3p3
S: 1s2, 2s2 2p6, 3s2 3p4

Question 13

Why is there a drop between P and S?

Answer 13

In P each of the 3p orbitals has 1 electron. In S one must have 2 electrons. The repulsion of these 2 electrons makes it easier to remove 1 of them.

Question 14

What is the difference between the metallic and non-metallic nature of the elements from left to right in period 3?

Answer 14

From left to right:
-the metallic nature decreases while the non-metallic nature increases.
-the tendency to gain electrons increases while the tendncy to lose electrons decreases

Question 15

Which elements of period 3 are metals, metalloids and non-metals?

Answer 15

Na, Mg, Al- Metals
Si- Metalloid
P, S, Cl, Ar- Non-Metals

Question 16

What are the structures of these elements?

Answer 16

Na,Mg, Al- Giant metallic
Si- Giant molecular
P, S, Cl- Simple molecular
Ar- Atomic

( Left to right:Giantsx2, simple, atomic)

Question 17

What types of bonding do these elements undergo?

Answer 17

Na, Mg, Al- Metallic
Si- Covalent
P, S, Cl, Ar- Van der waal forces

Question 18

What effect does the type of bonding play on the melting?

Answer 18

Metallic bonding:
- There are strong electrostatic forces of attraction between the positive ions and the negative delocalized electrons. Therefore lots of energy is required to break these bonds. Hence they have a high melting point. There is an increase in energy needed from Na—> Al because the ionic charge increases hence the F.O.A increases

Giant Covalent Lattice:
- Covalent bonds are strong and since they are many it takes alot of energy to break these bonds. Hence they have a high melting point.

Simple Molecular:
-The weak intermolecular forces between the molecules makes it easy to break the bonds. Less energy needed=Low melting point

Question 19

In the simple molecular elements how does size impact the melting point?

Answer 19

The larger the molecule the stronger the Van der waal forces are between the molecules. Therefore more energy would be needed to break the bonds of the molecules hence the melting point would increase.
Example Sulphur
Sulphur is larger than Phosphorus and therefore would need more enery despite it being further down in the period than P.

Question 20

What allows an element to be a great conductor?

Answer 20

It must have delocalised electrons in specifically and metallic stucture.

Question 21

Which elements are conductors and which is the best?

Answer 21

The metals; Na, Mg and Al are conductors as they contain delocalize electrons and have a metallic structure. Al is the best conductor as it has the highest number of delocalised electrons in its lattice.

Question 22

What is a semi-conductor and give an example?

Answer 22

This is an element which can conduct under certain conditions only eg Si.

Question 23

What are non-conductors?

Answer 23

These are elements which cannot conduct due to the absence of mobile, chagred particles. These elemetns include; P,S,Cl,Ar

Question 24

What is density?

Answer 24

Density is how much space an object takes up.

Question 25

Where in period 3 is there a general increase in density and why?

Answer 25

Between Na to Si there is an increase in density because the atoms get heavier and are pulled closer together due to increasing strength is bonds.

Question 26

Where is there a general decrease in density and why?

Answer 26

There is a drecrease in density from P to Ar because as the intermolecular forces get weaker, the molecules are further apart. (Think about gas vs solid which is more dense?)